摘要:

Equilibrium and Kinetic Acidities of Carbon Acids BY F. G. BORDWELL Northwestern University Evanston Illinois U.S.A. Received 1st May 1975 Bnomalous Bronsted coefficients observed for nitroalkanes are rationalized in terms of a mechanism involving an intermediate anion. The results of measurements of equilibrium acidities of carbon acids in dimethyl sulphoxide (DMSO) solution are presented and compared with similar measurements made in solvents of low dielectric constant and in water. A linear correlation between gas-phase acidities and DMSO solution acidities for a series of ketones and for a series of nitriles is presented. Equilibrium measurements indicate the presence of a high degree of strain in the anion derived from trifluoromethylsulphonylcyclopropane.This result shows that the sulphur atom in the trifluoromethylsulphonyl group is entering into strong conjugation with the carbanion. Nitroalkanes are the only monofunctional carbon acids that are acidic enough to allow equilibrium measurements to be made in protic media. As such their proton transfer reactions have been subject to extensive study. Interest in nitroalkane substrates has been intensified recently by the observation that the Bronsted a coefficient relating the rates of deprotonation by a given base and the equilibrium constants for a series of nitroalkanes is usually anomalous in that a is often larger than 1.0 and is occasionally less than zero (i.e. negative). This observation raises doubts as to the suitability of the common practice of using the size of the Bronsted coefficient as an index of the extent of proton transfer in the transition state for the deprotonation of a carbon acid.The observation also makes dubious the common practice of using rates of deprotonation (" kinetic acidities ") as an index of carbanion stability since this requires an assumption to be made concerning the size of a. A number of interpretations of anomalous Brijnsted coefficients have been given. Our own working hypothesis is that deprotonations of nitroalkanes are not simple one- step proton transfer reactions but instead proceed by way of an intermediate or a "virtual '' intermediate anion.4 In the present paper several examples of anomalous Bronsted coefficients will be given and interpreted on the basis of this two-anion mechanism.A method for measuring equilibrium acidities of weak acids in dimethyl sulphoxide (DMSO) solution will then be presented and discussed. (Such measure- ments when coupled with rate studies in DMSO will in time provide additional information concerning the mechanisms of proton transfers). Finally the use of kinetic acidities as a measure of the stabilities of carbanions will be examined. ANOMALOUS BRONSTED COEFFICIENTS AND THE TWO-ANION MECHANISM The first examples of anomalous Bronsted coefficients to which attention was drawn were deprotonations of a-arylnitroalkanes and a-methylnitroalkanes. It was observed that the Bronsted a obtained on deprotonation of a-arylnitroalkanes such as ArCH2N02 by a given base (with changing m-and p-substituents) was larger than 1.0 and that a in the series CH3N02 MeCH,NO, Me,CHNO, for deproton- ation by hydroxide ion was negative2 To these examples we can now add the deprotonation of nitroalkanes of the type GCH2CH,CH,N02 by a given base.100 F. G. BORDWELL Using 8 G substituents a Bronsted plot of log k for deprotonation by lyate ion in 50% MeOH-H,O against equilibrium acidities in 50% MeOH-H20 showed that as in the ArCH2N02series the rates were more sensitive to substituent effects than the equilibria (a = 1.56).5 It is possible to rationalize this result as well as the previous results on the basis of the following two-anion mechanistic scheme I B-+H-C-NO~ ,+ BI....H I I.slow ,fast 3 2 Geometric and solvent reorganization The first (reversible) step is formation of a weak H-bonded complex (1) a step somewhat similar to the first step in the Eigen mechanism for proton transfer.6 The second rate-limiting step is formation of the "essentially pyramidal "nitro carbanion 2.Some rehybridization and solvent reorganization must occur in this step but the major changes of this type are assumed to occur in the next (fast) step where the strong BH-.C- H-bond is broken and replaced by strong H-bonding to the oxygen atoms of the planar nitronate ion (3). According to this scheme structural changes on the rate will be determined primarily in the step where 1 is converted to 2 whereas structural changes on the equilibrium will be determined primarily by their effect on the stability of the planar nitronate ion 3.It is not unreasonable then to expect that a structural change may have one effect on the transition state leading to anion 2 and a dzferent (smaller greater or opposite) effect on the stability of anion 3. When the structural change causes a greater effect on the stability of anion 2 than anion 3 as is true for deprotonations of ArCH,NO or GCH2CH2CH2N02 systems a will be larger than 1.0. It is understandable that a change in Ar will have a large effect on the deprotonation rate in ArCH,NO since Ar is attached directly to the carbon atom bearing the charge in 2 and the transition state for deprotonation leading to 2 will be strongly influenced by changes in Ar. On the other hand in ArCH=NO (3) substituents will have a smaller effect since the negative charge is primarily on oxygen two atoms removed and is much less subject to the influence of Ar.A similar argument can be made for the GCH,CH,CH,NO system. It is reasonable then for kinetic acidities to be more sensitive to structural changes in Ar or G than are equilibrium acidities (a > 1.O). If structural changes cause opposite effects on the stabilities of anions 2 and 3 a negative a will result. For example there is good reason to believe that the increased equilibrium acidities with increased methyl substitution in the series CH3N02 MeCH,NO, Me,CHNO is caused by methyl stabilization of 3. It is entirely possible on the other hand that methyl substitutiou may destabilize anion 2 (e.g. by weakening the H-bond).The inverse Bronsted correlation can be rationalized in this way. The large (nearly maximum) kH/kDisotope effects observed for deprotonation of simple nitroalkanes such as nitromethane nitroethane etc. by hydroxide ion can ACIDITIES OF CARBON ACIDS also be rationalized nicely by the two-anion mechanism. In the overall sense these reactions are exoenergetic and the kH/kD isotope effect would be expected to be small. On the other hand if the conversion of 1 to 2 is considered to be rate limiting this conversion would not need to be exoenergetic and could have a near-maximum isotope effect. In discussing this mechanistic scheme the question has been raised as to whether or not the conversion of 2 to 3 requires an activation barrier.4 In this connection we may note that in a rather closely analogous reaction the deprotonation of 2,4,6-trinitrotoluene by EtO- in EtOH a careful analysis by Caldin indicates that solvent reorganization is the major factor contributing to the activation barrier.' Some solvent reorganization and bond rehybridization must occur in the conversion of 1 to 2 but it seems reasonable to suppose that an additional barrier must be surmounted in converting 2 to 3.The importance of the contributions from solvent and geometric reorganization to the activation barriers in proton transfers is indicated also by the work of Ritchie in DMSO solution.8 Solvent reorganization would be expected to be less in DMSO than in a protic solvent such as MeOH and indeed Ritchie finds ca.a 100-fold faster rate for a proton transfer in DMSO as compared to a reaction with a comparable change in free energy carried out in MeOH8 We anticipate that the two-anion mechanism for deprotonation will hold for other substrates where conversion to the final anion requires extensive geometric and solvent reorganization such as in the deprotonation of aldehydes and ketones. On the other hand for substrates where these factors are of less importance such as sulphones and nitriles a change in mechanism is indicated since these carbon acids behave "normally " in Bronsted correlations i.e. in a manner similar to oxygen and nitrogen acids. EQUILIBRIUM ACIDITIES IN DIMETHYL SULPHOXIDE (DMSO) If solvent reorganization presents less of an activation barrier in DMSO than in protic solvents as seems likely,8 the activation barrier between anions 2 and 3 will be less in DMSO and the mechanism of the deprotonation reaction should change.If the activation barrier disappears completely such nitroalkanes may give " normal " Bronsted correlations in DMSO. As the first step toward investigation of this matter we have examined the equilibrium acidities of nitroalkanes and many other types of TABLE1.-ABSOLUTEEQUILIBRIUM ACIDITIES FOR METHANE CARBON ACIDS CHSEWG IN DIMETHYL SULPHOXIDE SOLUTION name formula pKa nit romet hane CHJNO;! 17.2 methyl trifluoromethyl sulphone acetophenone acetone CH3 SOiCFj CH3COPh CH3COCH3 18.8 24.7 26.5 methyl phenyl sulphone dimethyl sulphone acetoni trile CH3S02Ph CH3SOZCH3 CH3CN 29.0 31.1 31.3 dimethyl sulphoxide CH3SOCH3 35.1 a not statistically corrected.carbon acids in DMSO solution. A method has been developed which allows absolute pK's reproducible to k0.05unit to be determined over a pK range from about 5 to 32.1° Acidity measurements on over 350 compounds have been made to date. Included in this survey are many of the common methane carbon acids of the type CH3EWG (EWG = Electron-Withdrawing Group; see table 1). Equilibrium F. G. BORDWELL data free of ion association effects have not been available hitherto for most of these weak acids. It will be observed that the acidities of CH3EWG range over 18 powers of ten; nitromethane is more acidic than the parent hydrocarbon methane by over 40 pK units.The major factor in determining the size of these enormous acidifying effects is believed to be the ability of EWG to delocalize the charge in the anion and thus to stabilize the anion i.e. CH,-EWG -+ CH,=EWG-. SOLVENT EFFECTS ON EQUILIBRIUM ACIDITIES OF CARBON ACIDS As might be anticipated from their wide range of acidities the acid strengths of carbon acids are highly solvent dependent. DMSO has a high dielectric constant (49 at 20°) a high dipole moment and a high degree of polarizability. As a con- sequence equilibrium acidities measured at low concentrations in this solvent are not complicated by ion association effects.1° Acidities in DMSO cannot be compared directly with those in solvents of low dielectric constant such as benzene ether or cyclohexylamine (CHA) since the latter are ion-pair acidities and cannot be placed on an absolute scale.When a series of hydrocarbons of similar structural types are compared such as fluorenes xanthenes triphenylrnethane diphenylmethane etc. relative acidities in DMSO and CHA are found to agree to within a pK unit or less. This agreement is fortuitous however apparently being a consequence of the extent of ion pairing in CHA remaining essentially constant throughout the series. This becomes evident once a carbon acid giving a localized anion such as phenylacetylene is included in the series. Now the apparent difference in ion pair acidities in CHA between fluorene and phenylacetylene is only 0.3 pK unit whereas the absolute difference as determined in DMSO is 5.9 pK units.Similarly in a (low dielectric constant) polyether solvent acetophenone appears to be more acidic than fluorene by ca. 4 pK units whereas absolute measurements in DMSO show that it is less acidic than fluorene by 2.3 pK units. This reversal is probably caused by ion pairing effects in the polyether solvent. Direct comparisons of absolute equilibrium acidities in DMSO and in water are possible (table 2). TABLE 2.-cOMPARISON OF EQUILIBRIUM ACIDITIES IN DMSO AND IN WATER acid pK (DMSO)a pK(H20)b benzoic 11.0 4.2 ni tromethane 17.2 10.2 acetone 26.5 20 9-cyanofluorene 8.3 11.2c malononitrile 11.1 11.0 bis(met hylsulphony1)met hane 15.O 12.7 a data from our laboratory unless otherwise noted. b ref (l) unless otherwise noted.C in 50 % MeOH + H20. Examination of table 2 shows that for acids in which the charge in the anion resides primarily on oxygen such as benzoic acid nitromethane and acetone the acidities are about 7 pK units greater in water than in DMSO. This is a consequence of the strong stabilization of oxide ions by H-bonding in water as compared to very weak stabilization by H-bonding in DMSO. On the other hand €€-bonding is a minor factor in stabilizing an anion in which the charge is delocalized over a carbon framework as in the anion derived from 9-cyanofluorene; here the acidity is some- what greater in DMSO than in water. Finally in acids where delocalization of the ACIDITIES OF CARBON ACIDS charge to the EWG function is relatively small such as malononitrile or bis(methy1- sulphonyl)methane the acidities in DMSO and water do not differ greatly.The stabilizing effect on an anion provided by solvation is enormous ranging from ca. 60-80 kcal/mol for the transformation from the gas phase to DMSO for ions such as ClO and C1-. It is not surprising then to find that the stabilizing effect of substituents is often far greater in the gas phase than in DMSO. For example p for equilibrium acidities of rn-andp-substituted benzoic acids is ca. 2.5 in DMSO,I1 as compared to ca. 10 in the gas phase.12 What is surprising is to find that examples exist where structural changes cause changes in DMSO solution acidities that approach those in the gas phase in size. Thus plots of DMSO solution acidities against gas- phase acidities l3 for a series of nitriles and for a series of ketones are linear with slopes of ca.1.1 and 1.3 respectively (fig. l).14 It would appear that when the 1210 14 '6> 18 22 26 30 34 38 42 46 50 equilibrium acidities in DMSO/kcal mol-' FIG.1.-Plot of gas-phase acidities (ref. (13) against DMSO solution acidities (ref. (14)). charge on the anion is highly delocalized the solvent effect remains essentially constant throughout the series and the effect of structural changes on the solution acidities gives a good measure of the relative intrinsic acidities of the compounds as revealed by gas-phase acidities. It is of interest to note that cyclopentadiene in which the charge on the anion is symmetrically distributed is more acidic in DMSO than in the gas phase judging from its position in fig.1 with respect to the lines for the nitrile and ketone families. THE EFFECT OF STRUCTURAL CHANGES ON EQUILIBRIUM AND KINETIC ACIDITIES For weak uncharged acids the effect of a structural change on the equilibrium acidity will generally give a good measure of the effect of structural change on anion stability. Since equilibrium acidity measurements for most types of carbon acids have not been available hitherto however it has become common practice to use kinetic acidities as an index of carbanion stabilities in order to obtain answers to F. G. BORDWELL such important questions as the effects of aromaticity antiaromaticity homoaro- maticity heteroatom substitution and s-character on relative carbanion stabilities.Some of the problems associated with the use of kinetic acidities as a means of judging carbanion stabilities have been outlined recently and several examples of misleading information relative to carbanion stabilities given by kinetic acidities with respect to a-hetero atom effects have been presented.l5 It appears likely that kinetic acidities have also given misleading information with respect to the effect of strain on the stability of substituted cyclopropane and cyclopropene anions.16 The kinetic acidities of such compounds have been used as an index of the presence of antiaromaticity in cyclopropene anions. For example the relative rates of base-initiated deuterium exchange for two phenylsulplionylcyclopropanes a phenylsulphonylcyclopropene and a corresponding open-chain compound are indicated under the formulas 4-7.Ph Ph Ph 4 5 6 7 Rel. rate 12 (1.0) Rel. rate 310 (1.0) The conclusion has been drawn that the faster rate for 4 relative to 5 indicates the absence of strain as a factor in determining the stability of the anion derived from 4 and that the lo3 slower rate for 7 compared to 6 can be attributed to antiaromaticity in the anion derived from 7.” Measurement of equilibrium acidities in DMSO for a number of cyclopropanes bearing EWG substituents i.e. c-PrEWG show however that the carbanions are highly strained relative to open-chain models (table 3). TABLE3.-EQUILIBRIUMACIDITIES OF CARBON ACIDS MEASURED IN DIMETHYL SULPHOXIDE SOLUTION EWG pK(CH3EWG) pK(i-PrEWG) pK(c-PrEWG) ApKa NO2 17.20+ 0.01 16.89+ 0.02 -27 -9 SOZCF 18.76+ 0.03 21.SO+ 0.03 26.60f0.03 7.3 COPh 24.70f 0.02 26.26+ 0.02 28.18+ 0.02 3.O SOzPh 29.04k0.05 >32 >32 >2.5 a pK(c-PrEWG)-pK(CH3EWG) statistically corrected.b rapid decomposition. Examination of table 3 shows that the c-PrEWG acids are less acidic than CH3EWG acids by factors ranging from ca. 3-9 pK units. This would correspond to an increased strain in the c-PrEWG- anion relative to the CH2EWG- anion amounting to 4 to 12 kcal/mol. The size of the estimated strain will depend of course on the open-chain model chosen. Use of the i-PrEWG model would decrease the size of the estimated strain but the conclusions would remain the same. In cyclopropane itself the high degree of s-character in the C-H bond is expected to cause an increase in acidity relative to open-chain alkanes.The EWG groups NO2 S02CF, CQPh and SO,Ph are all apparently imposing a high degree of p-character on the carbanion leading to large strain effects in the anion which completely over- shadow an intrinsically greater acidity caused by a higher degree of s-character in the C-H bond in the undissociated acid. The results show that strain effects in c-PrEWG- anions are large which indicates that strains in the corresponding cyclo- ACIDITIES OF CARBON ACIDS propene anions will also be large and must be considered in assessing their anti- aromaticity. The results also indicate a strong conjugation of sulphur in the S02CF3 and S02Ph groups with the carbanion the order being S02CF3> S0,Ph.This is the order predicted by Craig with respect to the change in the ability of d-orbitals on sulphur to participate in conjugation on introduction of a strong EWG (CF3 for Ph in the present instance). Recent theoretical analyses suggest however that d-orbital conjugation may not be significant.lg It is possible therefore that orbitals other than d-orbitals may be involved but there can be no escaping the fact that the conjugative ability of sulphur in sulphonyl groups towards a-carbanions is strong. Financial support by the National Science Foundation (MPS 74-12665) is grate- fully acknowledged. The author also wishes to express his appreciation to the University of Wales for a Visiting Professorship during the tenure of which this paper was written.R. P. Bell The Proton in Chemistry (Cornell University Press 2nd edn. 1973). ’F. G. Bordwell W. J. Boyle and J. A. Hautala J. Amer. Chem. Soc. 1969 91 4002. R. A. Marcus J. Amer. Chem. SOC., 1969,91,7224; A. J. Kresge J. Amer. Chem. Soc. 1970 92 3210. F. G. Bordwell and W. J. Boyle J. Amer. Chem. SOC., 1975 97,3447. F. G. Bordwell and J. E. Bartmess unpublished results. Ivl. Eigen Angew Chem. (Int. Ed) 1964 3 1. ’E. F. Caldin J. Chem. SOC., 1959 3345. C. D. Ritchie J. Amer. Chem. SOC.,1969 91 6749. R. P. Bell and B. G. Cox J. Chem. SOC.B 1971 654; F. Hibbert and F. A. Long J. Amer. Chem. SOC.,1971 93,2836; 1972 94,2647. lo W. S. Matthews J. E. Bares J. E. Bartmess F.G. Bordwell F. J. Cornforth G. E. Drucker Z. Margolin R. J. McCallum G. J. McCollum and N. R. Vanier J. Amer. Chem. SOC.,1975 97 Oct. issue. I1 1. M. Kolthoff and M. K. Chantooni J. Amer. Chem. SOC.,1971,93,3843 ; C. D. Ritchie and R. E. Uschold J. Amer. Chem. Soc. 1968,90 2821. R. Yamdagni T. B. McMahon and P. Kebarle J. Amer. Chem. Soc. 1974,96,4035. l3 T. B. McMahon and P. Kebarle J. Amer. Chem. SOC., 1974,96 5940. l4 F. G. Bordwell J. E. Bartmess G. E. Drucker W. S. Matthews and Z. Margolin J. Amer. Chem. SOC., 1975,97,3226. l5 F. G. Bordwell W. S. Matthews and N. R. Vanier J. Amer. Chem. SOC.,1975 97,442. l6 F. G. Bordwell N. R. Vanier W. S. Matthews W. B. Hendrickson and P. W. Skipper J. Amer. Chem. SOC.,1976. l7 R. Breslow Acc. Chem. Res. 1973 6 393. D. P. Craig and E. A. Magnusson J. Chem. Soc. 1956,4895. l9 S. Wolfe A. Rauk and I. C. Csimadia J. Amer. Chem. Soc. 1969 91 1567; J. I. Musher J. Amer. Chem. SOC., 1972 95 1370.

ISSN:0301-5696    

DOI:10.1039/FS9751000100

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Kinetic and Thermodynamic Basicities of Anions in Mixed Solvents BY BRIAN G. Cox AND ALANGIBSON Department of Chemistry University of Stirling Stirling Scotland Received 1st May 1975 The rates and equilibria for reactions involving proton transfer between several nitroalkanes and fluoride acetate and hydroxide ions have been investigated in dimethylsulphoxide+ water and trifluoroethanol+ water mixtures. The results show that the effects of solvent variation on the rates are not simply related to the effects on the overall free energy changes occurring during the reactions. Recently observed anomalies in Bronsted /3 coefficients for reactions involving these substrates disappear as the fraction of water in the dimethylsulphoxide + water mixtures decreases. It is argued that the anomalous values in water result from specific hydration effects on the acidity constants.Much of the recent interest in proton transfer reactions (eqn (1)) SH+B-+ S-+BH (1) has centred on the relation between the reaction rates and kinetic hydrogen isotope effects and the overall free energy change occurring during the reacti0n.l. The effect of variations in the nature of both SH and B- have been extensively studied generally in aqueous media. It is becoming increasingly clear however that acid strengths are considerably influenced by the strong solvation which occurs in water. It is noticeable for example that within a series of related acids (e.g. benzoic acids carboxylic acids phenols) the pK,’s in dipolar aprotic solvents (polar solvents without acidic hydrogens) and in the gas phase are considerably more sensitive to the effect of substituents than they are in water.It has been suggested that a number of apparent anomalies in Bronsted p coefficients relating rates and equilibrium constants for proton transfer reactions may in fact have their origin in the variable contribution of hydration to the acidity constants (of SH and BH)? In the present paper we report a study of the effect of changing solvent on the rates equilibria and Bronsted p coefficient for reactions of the type shown in eqn (I). The substrates used include several substituted nitroalkanes and 1,l-dinitroethane. These substrates are sufficiently acidic to enable their pK,’s to be determined accur- ately.Also reactions involving these substrates in aqueous solution have in many cases shown anomalous Bronsted /3 coefficient^.^. It is of interest therefore to see the extent to which these anomalies are due to the effects of solvation in water. EXPERIMENTAL MATERIALS 2-Nitropropane was purified by spinning band distillation. G.l .c. and n.m.r. analysis showed it to be free from traces of 1-nitropropane and nitroethane. Nitroethane was dried (MgS04) and distilled. 1,l-Dinitroethane was prepared by reaction of nitroethane with silver nitrite in basic solution as previously described.* 1-Arylnitroethanes 1-(3’-107 ANION BASICITIES IN MIXED SOLVENTS methoxypheny1)-and 1-(4'-nitrophenyl)-l-nitroethane were prepared by oxidation of the oximes of the corresponding acetophenones as described by Bordwell et aIa9 Dimethylsulphoxide was purified by distillation under reduced pressure from calcium hydride ; trifluoroethanol was dried (Na2S04) and distilled ; inorganic chemicals were of Analar Grade.pK MEASUREMENTS pK,'s of the various acids in Me2SO+H20 and CF,CHzOH+H20 mixtures were determined by potentiometric measurements with glass electrodes. The electrodes were calibrated with dilute (ca. 10-'-10-3 M) perchloric acid solutions in the solvent mixtures. Values for acetic acid agree within 0.1 pK unit with values previously determined using conductance measurernents,l0 acid-base indicators and potentiometric measurements with hydrogen e1ectrodes.l2 In the determination of the pK of HF allowance for the formation of HFj was made using simultaneous measurements of pH and the fluoride ion activity (with a fluoride-ion-selective electrode) as described by Kresge and Chiang.l3 The pK of rn-methoxy phenylnitroethane in water was obtained from spectrophotometric measurements in tris( hydroxyme t hyl amino)met hane buffers. The acidity constants KA of the acids HA in solvent S are represented by eqn (2) where the y's are activity coefficients referred to infinite dilution in solvent S. The activity coefficients of ionic species were calculated from the Davies equation (3) l4 where A is the Debye-Huckel function dependent upon the solvent dielectric constant and temperature. The activity coefficients of AI+ -logy* =-0.3AI (3) l+P neutral species were assumed to be unity.All measurements were carried out at 25 (+0.1)"C. KINETIC MEASUREMENTS The proton-transfer reactions were followed either by direct spectrophotometric obser- vation of the substrate anion formed or by measuring spectrophotometrically the rates of iodination of the substrates. Iodination reactions were followed under conditions where the rate determining step in the iodination was the deprotonation of the substrate (SH). Thc particular method used for a given substrate-base reaction is included in the Results section. The reactions were followed with a Gilford 2400 spectrophotometer or a Durrum-Gibson stopped flow spectrophotometer depending upon the reaction rate. All measure- ments were made at 25 (kO.2)"C. RESULTS 2-NI TROPROPAN E The rates of deprotonation of 2-nitropropane (2-NP) were obtained from measured rates of iodination.15* l6 The reactions were carried out in solutions with [2-NP]ca.5 x 10-3-5x 10-2M [I-] = 0.01 M and initial iodine concentrations <10-4M. The rates were measured in Me,SO +H20 mixtures using acetate and fluoride buffers with anion concentrations in the range 2 x 10-2-10-'M. Fluoride buffers with [F-]/[HF] 21 40 (pH = 4.7 in water) were used to avoid complications arising from HF; formation. The iodination of 2-nitropropane is known to be reversible 16. l7 where either [I-] or [H+]is high but under the conditions used the rate law shown in eqn (4) and (5)was obeyed over -d[I;]/dt = -d[2-NP],dt = ke[2-NP] (4) B. G. COX AND A. GIBSON 109 where k = k,+k~[B] (5) at least 90 % of reaction.acetate or fluoride ions. In eqn (4) and (9,[I;] = [I2]+[I,] and B refers to either In solvent mixtures containing dimethylsulphoxide ko was negligible compared with kB[B] so that k = kB[B]. The results for the system are given in table 1. NITROETHANE Rates of proton transfer from nitroethane (NE) to acetate in Me2S0+H20 and CF3CH20H+H20 mixtures have been measured. Equilibrium constants have been obtained from measured pKa's. The experimental conditions and methods were the same as for 2-nitropropane except that nitroethane concentrations were ca. 0.15 M. The results are also listed in table 1. TABLEN RATES AND EQUILIBRIA FOR PROTON TRANSFER FROM NITROALKANES (SH) TO ACETATE AND FLUORIDE IONS IN SOLVENT MISTURES AT 25°C 1.2-Nitropropane in Me,SO+ HzO mixtures -PKHOA~6 10gkoAcb ~KSH XM~~SO~ ~KSH ~KSII -PKHF 7+hgFb 0.00 7.74c 3.0 0.43 4.6 0.42 0.21 9.5 3.7 1.41 4.6 1.95 0.38 11.0 3.9 2.25 4.3 3.16 0.51 12.1 4.1 2.74 3.9 3.89 0.70 -3.80 1.oo 16.2d 4.2 2. Nitroethane in Me2SO+Hz0 and CF3CH20H+ HzO mixtures Xnre2soa ~KSHPKsIi-pKHoAc 5+bgkOAcb XTFE' PKSH ~KSH -~KHOAC 5+logkOAC 0.00 8.8 4.1 0.51 0.00 8.8 4.1 0.51 0.21 10.2 4.2 1.63 0.20 10.2 4.4 0.22 0.38 11.7 4.5 2.47 0.32 -0.17 1.00 16.4d 4.4 -0.50 11.5 4.4 -0.07 (a) mole fraction of dimethylsulphoxide ; (6) kg dm3m01-'s-' are rate constants for proton transfer from SH to B ; (c) ref. (16) ; (d)ref. (19) ; (e)mole fraction of trifluoroethanol.TABLE 2.-cOMPARISON OF pKa'S AND RATES WITH ACETATE IONS FOR NITROETHANE AND 1,l-DINITROETHANE IN MezSO+HzO MIXTURES AT 25°C -l,t dinitroethane nitroethane 7-XMe2SOQ PKSH 5 fIogkOAc PKSH 5+lOgkoAcb 8' 0.00 5.24d 5.33d 8.8 0.51 1.34 0.21 3.3 7.15 10.2 1.63 1.12 0.38 5.3 8.55 11.7 2.47 0.95 1.oo 6.6' - 16.4' - - (a)mole fraction of Me,SO ; (6) ko~~/dm~mol-'s-' are rate constants for proton transfer from SH to QAc- ; (c) = [~O~~~~,(DNE)-~O~~~A,(NE)]/[~K(NE)-~K(DNE)] ; (d) ref. (7) ; (e) ref. (19). 1,l-DIN ITROET H A N E The reaction of 1,1 -dinitroethane (DNE) with acetate in Me2S0+H20 mixtures (eqn 6)) was followed by observing the appearance of the dinitroethane anion koAc-CH3* CH(N02)2 +OAC-+ CH,* C(NO2); +HOAC (6) kHoAc ANION BASICITIES IN MIXED SOLVENTS spectrophotometrically at 385 nm.DNE concentrations were ca. 5 x M and acetate concentrations were in the range 0.02 to 0.1 M with [OAc-] = 2[HOAc]. The observed rate law was of the form shown in eqn (7) and (8) where K = KDNE/KHOAc is the equilibrium constant for -d[DNE]/dt = k,[DNE] where k = kOAc[OAc-]{ 1 + [HOAc]/K,[OAc-]) reaction (6) and koAcis the rate constant for the forward reaction. Kinetic and thermodynamic data for the reaction between dinitroethane and acetate are given in table 2 together with the corresponding results for nitroethane. 1-ARY LNI TROETH A NES The rates of reaction of m-OMe- and p-NO2-phenylnitroethane with OH-were followed by spectrophotometric observation of the nitroalkane anion.* Nitro-alkane concentrations were < M and [OH-] ca. 10-3-2x M. The observed rate law was as shown in eqn (9) -d[nitroalkane]/dt = ko,[OH-][nitroalkane]. (9) TABLE 3.-pK,'s AND RATES WITH HYDROXIDE AND ACETATE IONS FOR 1-ARYLNITROETHANES YC6H4/CH(Me)N02,IN Me2SO+H20 MIXTURES AT 25°C (1) reaction with OH--Y = pNO2 Y = m-OMe xble2SOa 'pKsHfIb p&H logkOHb B" 0.00 6.63 1.70 7.40 0.93 1.12 0.21 7.40 3.15 8.93 1.93 0.80 0.51 8.64 5.13 10.94 3.63 0.64 (2) reactions with OAc-Y = P-NO~ Y = m-OMe -xMe2s0 ~KSH 10gkoAcb -0 PC 0.51 8.64 0.20 10.94 -1.38 0.69 (a) mole fraction of Me2S0 ; (b) kg/drn3mol-'s-' are rate constants for proton transfer from SH to B ; (c) /3 = [b&B (p-NO2)-logk~(rn-OMe)]/[pK(m-OMe)-pK(p-NO,)].The rates of reaction with acetate ions in 80 vol% Me2SO+H,0 were also measured the reaction of the p-NOz derivative being followed by observation of the nitroalkane anion and the rn-OMe derivative by iodination as described earlier. Kinetic and thermodynamic data for the reactions are listed in table 3. DISCUSSION For a proton transfer reaction such as that shown in eqn (l) the Bronsted fl coefficient relating the effect of variations in the nature of B-or SH on the rates and equilibria may be defined by eqn (lo) where k is the p = 6AG*/6AG0 = GlogkB/610gK (10) rate constant and K (= KSH/KBH) is the equilibrium constant for the reaction. A Bronsted exponent greater than unity then indicates a reaction in which the substituent B.G. COX AND A. GIBSON effect on AG* is greater than that on AGO which is contrary to the expectation that the effect of a substituent should vary monotonically with the extent of reaction. Examination of the results in table 1 shows that for the reactions studied the effects of solvent on the reactions do not vary monotonically with the extent of proton transfer during the reaction. The pK,’s of both the substrates and bases vary con- siderably with increasing amounts of Me,SO (or CF3CH20H) in the solvent but there is little change in the equilibrium constant for the reactions between the substrate and bases (expressed as pKsH -pKBH). The rates however increase rapidly with content of the organic component in Me,SO+H,O mixtures and decrease in CF,CH20H+H,0 mixtures.This behaviour can be readily explained in terms of the large difference in the interaction of water with anions of high charge density such as F- RCO, RNO; (capable of H-bond stabilisation) and large polarisable anions (such as the transition state anions) relative to dipolar aprotic solvents such as Me2S0.18 Trifluoroethanol being more acidic than water would be expected to stabilise even further the smaller anions relative to the transition state anions. Such behaviour is perhaps not unexpected as it has been known for some time that the rates of reactions shown in eqn (1 1) *X-+RX + RX* +X-(1 1) where X and *Xare isotopes are very sensitive to the nature of the solvent IS although the equilibrium constant for the reaction is clearly independent of the solvent.reaction coordinate FIG.1.-Bronsted coefficients and solvent effects. The relevance of these effects to the explanation of the origin of p values greater than unity in water can be seen from the results in tables 2 and 3. Considering for example the results in table 2 as the fraction of Me,SO in the solvent is increased the difference in pK of dinitroethane and nitroethane increases going from 3.6 pK units in water to 9.8 pK units in Me2S0.19 This reflects the loss in the higher hydration energy of the mononitroalkane anion relative to the dinitroalkane anion with its more highly dispersed charge. Similar effects do not operate in the transition states since for both reactions these have highly dispersed charges.Thus the B ANION BASICITIES IN MIXED SOLVENTS values decrease (and become normal) as the Me2S0 content of the solvent increases. The results in table 3 may be explained in an analogous manner. This explanation is shown schematically in the figure where curve I represents nitroethane (or nz-OMe-phenylnitroethane) and curve I1 represents dinitroethane (or p-NO2-phenylnitroethane). The full lines indicate the reaction in for example the gas phase and the dotted line in water. For simplicity the curves in the two phases have been moved relative to one another to align reactants and the solvent is shown as having no effect on the transition state. It is interesting to note that Jones and Patel 2o have recently found that comparison of the effects of increasing fluorine substitution in acetylacetones on the acidity and rates of detritiation leads to anomalous p values.This is attributed to the effects of hydrate formation of the fluorinated ketones on their measured pK,'s. This may be regarded as an extreme example of the effects of hydration operating specifically on the measured equilibrium constant and so not having a corresponding effect on the transition state energies. It has in fact been suggested many years ago 21 that hydration effects could simply explain " anonlalies " such as the observation of parallel but separate Bronsted plots for primary secondary and tertiary amines. We suggest that in addition because within a series of related acids increases in acidity caused by substituents frequently arise from increased delocalisation of the anion charge hydration can also alter the slopes of Bronsted plots.This will apply equally to results obtained from the variation of a series of substituted carboxylate anions with a constant substrate or from variations in the nature of the substrate with a constant base. Such a factor should be borne in mind when attempting to relate observed fl values to transition state structures etc. We thank the SRC for a studentship for A. G. and a research grant and Prof. R. P. Bell for discussion. R. P. Bell Chenr. SOC.Reu. 1974 3 513. A. J. Kresge Chem. SOC. Rec.. 1973 2 475. (a) I. M. Kolthoff and M. K. Chantooni J. Amer. Chem. SOC.,1971 93 3843 ; (6)B.G. COX Ann. Reports Chem. SOC. A 1973 249. K. Hiraska R. Yamdagni and P. Kebarle J. Amer. Chent. SOC.,1973 95 6833. e.g. F. G. Bordwell and W. J. Boyle J. Amer. Chem. SOC.,1972 94 3907. B. G. Cox and A. Gibson Chein. Comnr. 1974 683. R. P. Bell and R. L. Tranter Proc. Roy. SOC. A 1974 337 517. 'R.B. Kaplan and H. Shechter J. Amer. Chem. Soc. 1961 83 3535. F. G. Bordwell W. J. Boyle and K. C. Yee J. Amer. Chem. SOC. 1970 92 5926. lo J.-P. Morel Brill. SOC. Chint. France 1967 1405. E. C. Baughman and M. M. Kreevoy J. Phys. Chent. 1974 78,421. l2 J.-C. HallC R. Gaboriaud and R. Schall Bull. SOC.Chint. France 1970 2047. l3 A. J. Kresge and Y. Chiang J. Phys. Cheni. 1973. 77 822. l4 C. W. Davies Zon Association (Butterworth London 1962) eqn (3.14). l5 R. P. Bell and D. M. Goodall Proc. Roy. SOC.A 1966 294,273. I' M. H. Davies J. C. S. Perkiri Ii 1974 1018. l9 R. P. Bell and E. Gelles Proc. Roy. SOC.A 1952 210 310. l7 A. J. Parker Chcm. Rev.,1969 69 1. V. M. Belikov Yu. N. Belokon and N. G. Faleev Izt'est. Akad. Narik S.S.S.R. Ser. Khim. 1971 2 272. 'O J. R. Jones and S. P. Patel J. Amer. Chem. SOC. 1974 96 574. 21 R. P. Bell J. Phys. Chem. 1951 55 885.

ISSN:0301-5696    

DOI:10.1039/FS9751000107

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Carbon-13 Isotope Effects on Proton Transfers from Carbon BY JOHN BANGER JAFFE AN-CHUNG ANNETTE LINAND WILLIAM JR.* H. SAUNDERS Department of Chemistry University of Rochester Rochester New York 14627 U.S.A. Receiiied 22nd April 1975 Carbon-13 isotope effects at C-2 have been determined on the rates of elimination from 2-phenylethyl-dimethylsulphoniumand -trimethylammonium salts with hydroxide ion in mixtures of dimethyl sulphoxide and water. The effects are consistently larger than those predicted from model calculations for a semi-classical (without tunnelling) isotope effect. The most reasonable inter- pretation is that the form of the dependence of the isotope effect on solvent composition is controlled by the semiclassical effect with a substantial tunnel effect superimposed.For some time deuterium and tritium isotope cffects ir? hydrogen-transfer reactions have been interpreted in terms of the Melander-Westheinier model,' * which predicts a maximum in the isotope effect when the hydrogen is half transferred. Bell and Goodall predicted that proton-transfer reactions should show this inaximgm when the pK values for the substrate and the conjugate acid of the attacking base were equal and presented evidence that this prediction was at least approximately true. Subsequently maxima in kH/kDvalues have also been achieved with sinzle substrates by varying the compositions of mixtures of aqueous hydroxide and dimethyl ~ulphoxide.~-~ Although this picture can be modified somewhat by tunnel effect^,^ non-linear transition states or inclusion of the proton transfer in a more complex reaction coordinate such as that for bimolecular eliii~ination,~.'~ it has stood without sub- stantial challenge until recently.Bordwell and Boyle argue that the maximuin in a plot of kH/kD against ApK is so broad and diffuse as to be of little value in determining transition-state syinmetry.ll Bell Sachs and Tranter conclude from calculations on an electrostatic model of the proton-transfer transition state that there is little variation in the stretching force constants of the base-proton and substrate-proton bonds as the base strength is varied and that observed kH/kDmaxima arise primarily from variations in the tunnel effect contribution. l2-I3 Rather large variations in the relative values of these force constants are needed to obtain isotope effects substantially below the maximum values from a semiclassical (without tunnelling) model.In view of the uncertainty currently surrounding the interpretation of hydrogen- transfer isotope effects we decided to undertake a study of carbon isotope effects on proton transfers from carbon. The reactions chosen to begin our study were the E2 reactions of 2-phenylethyl-dimethylsulphonium and -trimethylainmonium salts with hydroxide ion in mixtures of water and dimethyl sulphoxide. Both had been shown to give k,/k maxima.5* The dependence of the semiclassical isotope effect lo and the tunnel effect l6 for 13C as compared with 12C had been calculated as a function of the extent of proton transfer by the Wolfsberg-Stern-Schachischneider programs.' 7*1 These two contributions to the overall isotope effect appeared to differ sufficiently to permit a clear assessment of their relative importaiice from experimental data.113 CARBON-13 ISOTOPE EFFECTS EXPERIMENTAL MATERIALS 2-Phenylethyldimethy1sulphoniumbromide and 2-phenylethyltrimethylammoniumbro-mide were prepared as previously described. 19-2 Inorganic chemicals and the solvents used for extraction were analytical reagent grade. Dimethyl sulphoxide was purified by refluxing over calcium hydride followed by fractional distillation. Mixtures of water and dimethyl sulphoxide were prepared by weight. REACTIONS Reaction mixtures consisted of 50-75 ml of solution 0.1-0.2 M in substrate and 0.1-0.5 M in sodium hydroxide.The solutions were prepared and the reactions allowed to run in a constant-temperature bath controlled to 0.05"C. Two techniques were used. For the slower reactions an excess of base was employed the reaction allowed to run to the desired extent of completion as estimated from available rate con~tants,~~~ the reaction mixture cooled in ice water and the extent of reaction confirmed by titration of unconsumed base. For fast reactions an insufficiency of base was used so as to stop the reaction at the desired extent of completion and the mixture checked for any remaining base by titration. Titration of reaction mixtures containing 2-phenylethyltrimethylammoniumbromide used thymol- phthalein as indicator with added pyridine to repress interference from the tertiary amine present.2*2 The neutralized reaction mixtures were extracted with five 30-ml portions of benzene to remove styrene and then evaporated to dryness on a rotary evaporator with a mechanical pump. For reaction mixtures containing high proportions of dimethyl sulphoxide water was added to the residue and the extraction and evaporation process repeated. Tests with simulated reaction mixtures showed that the styrene was completely removed by this procedure. DEGRADATION OF UNREACTED SUBSTRATE The residue obtained above was dissolved in 100 ml of water. About 1.5 g of sodium carbonate was added followed by the cautious portionwise addition with stirring of potassium permanganate until its colour persisted (at least 10 g).After the reaction subsided the mixture was heated on a water bath for 3 h cooled in an ice bath and cautiously treated with sodium sulphite to destroy excess permanganate followed by acidification (litmus) and stirring for an additional half hour. The mixture was then extracted with three 100-ml portions of ether the extracts dried over magnesium sulphate and the ether removed. The residue was recrystallized from hot water to give benzoic acid in essentially quantitative yield. Its purity was checked spectroscopically on selected samples. A 30-40-mg sample of benzoic acid was dissolved in 6-9 ml of concentrated sulphuric acid in a 50-ml 3-neck flask. Dry nitrogen was bubbled through the flask and an attached collection trap for 30 min.The trap was then immersed in liquid nitrogen the flask heated to 40-60",and a side arm containing 20-30 mg of sodium azide rotated so as to add the azide to the flask. The reaction was allowed to proceed for one hour. The collection trap (still cooled by liquid nitrogen) was removed from the flask attached to a vacuum line pumped on carefully to remove nitrogen and then evacuated for 15 min with a mercury diffusion pump. It was then transferred to a gas chromatograph inlet and the contents swept by a helium flow of 25 ml/min onto a 30-ftx 1/4-in column of 20% adiponitrile on Chromosorb P. The carbon dioxide was collected in a trap cooled by liquid nitrogen. The trap was transferred to the vacuum line and the carbon dioxide condensed in a drying trap containing phosphorus pentoxide.The carbon dioxide was then transferred to a mass spectrometer sample tube. ISOTOPE RATIO DETERMINATION Sample tubes containing the carbon dioxide from the original substrate and from substrate recovered after partial reaction were attached to the dual viscous inlet system ofan J. BANGER A. JAFFEE A,-C. LIN AND w. H. SAUNDERS JR. 115 Atlas CH-4mass spectrometer equipped with dual Faraday cup collectors and the ratio of the m/e 44/45 values determined for each sample. There was some dependence of apparent isotope ratio on sample size. In the work with the sulphonium salt a plot of apparent isotope ratio against sample size (measured as total ion current) was prepared using tank carbon dioxide and was used to correct the sample isotope ratio to the value it would have at the same pressure as the standard.Improved precision was obtained in the work on the ammonium salt by keeping sample and standard sample total ion currents within f10% of each other and running the standard both before and after each sample. Observed 44/45 ratios were corrected for the natural abundance of 170 in the sample. CONTROL EXPERIMENTS Carbon isotope ratios (44145) from three samples of carbon dioxide from the same tank were the same to within rfrO.l%. Five samples of the same lot of benzoic acid were decarboxylated to give carbon dioxide samples of the same 44/45 ratio to within +0.2%. Two identical samples of 2-phenylethyldimethylsulphoniumbromide were dissolved in base- solvent mixtures containing pure water and 40% dimethyl sulphoxide respectively and recovered before any significant reaction could occur.They were degraded to give carbon dioxide of the same isotope ratio to within +0.2%. Another three samples of the sulphonium salt were degraded (without prior dissolution) to give carbon dioxide of some- what more variable isotope ratio (+0.4%) but the sample sizes as measured by total ion current varied over a 30%range. Finally there was no measurable attack of base on the glass reaction flasks over the time of the longest reactions. RESULTS The substrates Ia and Ib have been previously shown to react with PhCH2CH2X Ia X = -SMe2Br Ib X = -NMe,Br hydroxide in mixtures of dimethyl sulphoxide and water to give quantitative yields of styrene.Samples of each substrate were oxidized with potassium permanganate 9 0 00 0 20 40 60 80 100 mole % water FIG.1.-Observed j3-12C/13C 1)x 100 for the reactions of isotope effects expressed as (kI2/kl3-2-phenylethyl-dimethylsulphoniumion at 30°,open circles and -trimethylammoniom ion at 60°C closed circles with hydroxide ion in mixtures of water and dimethyl sulphoxide. to give benzoic acid which in turn was decarboxylated by treatment with hydrazoic and sulphuric acids to give carbon dioxide whose m/e 44/45 ratio was determined on an isotope-ratio mass spectrometer. Control experiments showed that the degradation was quantitative and gave isotope ratios reproducible to & 0.2 %. The elimination CARBON-13 ISOTOPE EFFECTS reactions were allowed to proceed to a specified extent and unreacted substrate isolated and degraded as above.The isotope ratio for the original sample (R"), the recovered sample (R),and the fraction of reactant remaining (F)were then substituted into eqn (1) to calculate the isotope effects. l4 The resulting values are collected k1&3 = log F/log(RF/R") (1) in table 1 and shown graphically as a function of solvent composition in fig. 1. Before the experimental results can be discussed the calculated values of the isotope effects must be considered. The transition-state model (11) was closely similar to that used earlier.g. lo The " extent of hydrogen transfer " was defined as the fractional weakening of the C-H bond.The extent of C-S bond weakening TABLE DIM CARBON ISOTOPE EFFECT FOR THE REACTION OF PhCH2CH2X WITH HYDROXIDE ION IN MIXTURES OF WATER AND DIMETHYL SULPHOXIDE X percent DMSOa ternp/'C* no. of detns. (klz/k13-l) x 100" SMe2 SMe SMe2 SMe SMe SMe §Me2 SMe SMe2 NMe NMe NMe NMe NMe NMe NMe NMe NMe 0.0 1.25 2.5 5.0 5.O 7.25 10.0 40.0 6.0 0.0 0.0 5.0 10.0 10.0 20.0 40.0 40.0 60.0 30 30 30 30 30 30 30 30 30 60 80 60 60 80 60 60 80 60 5 5 6 8 3 7 10 7 3 3 4 4 4 3 3 3 4 4 2.17$-0.18(0.50) 2.23f0.S4(0.39) 3.05 0.17(0.44) 3.40+ 0.14(0.33)d 3.15+ 0.08(0.34)e 1.925 O.OS(0.12) 2.32+ 0.09(0.20) 2.01 +0.09(0.22) 1.83+O.l l(0.47) 2.13 f0.05(0.22) 1.54+0.13(0.41) 1.64f O.Og(0.29) 1.61f0.04(0.13) 1.4650.03(0.13) 1.57+0.05(0.22) 1.61f0.14(0.60) 1.69f 0.13(0.41) 2.105 0.15(0.48) a mole percent ; b -t 0.1"C; C percent isotope effect.The figure immediately following the & is the standard deviation of the mean while the figure in parentheses is the 95% confidence limit. d results obtained by A. Jaffe ; e results obtained by A.-C. Lin under same conditions as for d. the extent of 0-H bond formation and the extent of C-C double bond formation were changed parallel to the extent of hydrogen transfer. A moderately curved potential barrier giving a maximum reaction coordinate frequency of ca. 9OOi cm-l was used.lg Other details of the model are described el~ewhere.~ lo The calculated isotope effects as functions of the extent of hydrogen transfer are given in fig. 2 for b-hydrogen as compared with deuterium and in fig. 3 for P-"C as compared with 13C.The figures also include the tunnel corrections calculated from the simple form of the Bell equation,13 which is valid for all values of the reaction-coordinate frequency of our model. This model should be valid for discussions of isotope effects for both Ia and Ib for the deuterium isotope effect is independent and the carbon isotope effect nearly independent of the nature of the leaving group.1° The effects of varying numerous parameters for I1 and related models have been explored.1° While numerical values can be changed by such manipulation the basic bell shape of the deuterium isotope effect curve and the S-shape of the carbon isotope effect curve remain essentially the same. J. BANGER A. JAFFE A.-C. LIN AND W. H. SAUNDERS JR.117 CH2-CH2 H H-0 I1 Intelligent comparison of the experimental results and the model calculations requires some knowledge of the relation between the solvent compositions used in the experiments and the " extent of hydrogen transfer ". No rigorous connection is possible but one can make the reasonable assumption that increasing dimethyl sulphoxide concentration leads to a decrease in the extent of proton transfer in the transition state because the hydroxide ion is becoming a stronger ba~e.~~'~' Second one can compare the experimental k,/k values (fig. 4) at the maxima with the smallest values on either side of the maxima. One then transfers these (kH/kD)max/(kH/kD)mi" values to the calculated kH/kDcurves of fig. 2. From such comparisons one can conclude that the pure-water result for the sulphonium salt represents an extent of hydrogen transfer of about 0.5 or a little more while 60% dimethyl sulphoxide represents about 0.3 extent of hydrogen transfer.Corresponding figures for the ammonium salt results (allowing for the temperature difference) are about 0.6 and 0.2 respectively. Whether or not the tunnel effect is included in the calculated values has only a minor influence on these ranges. 01 03 05 07 09 extent of H-transfer FIG.2.-Calculated values of k~/k~ for elimination from ethyldimethylsulphonium ion at 25°C. See text and ref. (9) and (10) for details of model. Closed circles semiclassical isotope effect ; open circles tunnel effects ; half-open circles combined semiclassical and tunnel effects.DISCUSSION The results for the sulphonium salt (Ia) in fig. 1 show a scatter at low proportions of dimethyl sulphoxide which considerably exceeds the combined standard deviations. The t-test 28 demonstrates that the results at 95.0 and 97.5% water are significantly higher with >99 % probability than all other results. Similar isotope effects (3.40 and 3.15 %) were obtained by two different workers at 95.0 % water. We can see no good reason however why the results at 100 and 98.75 % water should be so different from those at 97.5 and 95.0% water. Some systematic error perhaps arising from the slowness of the reactions in the highly aqueous media may be involved. We CARBON-23 ISOTOPE EFFECTS assume that the results follow a smooth curve of the sort depicted in fig.1 giving approximately equal weight to all points in the 95-100% water region. Even if one set should be incorrect only the slope not the basic shape of the curve would differ. The shape of the dependence fits best either the semi-classical isotope effect or the isotope effect with tunnelling of fig. 3. It might also fit though less well the tunnel effect curve over the restricted range of extent of hydrogen transfer (ca. 0.3-.05) covered by the results. A pure semiclassical isotope effect can be excluded by the magnitude of the observed effects. The lowest observed value is 1.8 % rather than the inverse effects predicted by the pure semiclassical model. Varying the parameters of the model can give larger numerical values than those depicted in fig.3,1° but no realistic set of parameters gives semiclassical effects as large as those observed in the 0.3-0.5 extent of hydrogen transfer region. The most reasonable interpretation of the results is that the shape of the dependence follows approximately the semiclassical isotope effect but with a tunnel correction somewhat larger (ca. 2.0-2.5%) than that in fig. 3 superimposed. A dependence dominated by the tunnel effect cannot be completely excluded however for the sulphonium-salt results alone. 0.1 0.3 0.5 0.7 0.9 extent of H-transfer Fro. 3.-Calculated ,9-1*C/'3C isotope effects expressed as (klZ/kI3-1) x 100 for elimination from ethyldimethylsulphonium ion at 25°C. See legend to fig. 2 for explanation of symbols.The results on the ammonium salt (Ib) provide a better test since they are more precise and cover a wider range (ca. 0.2-0.6) of extents of hydrogen transfer. Although they describe a relatively shallow curve as the solvent composition varies the t-test 22 gives a >99 % probability that the value in 80% water is lower than the value in 100%water and >95% probability that it is lower than the value in 40% water. The ammonium-salt data are clearly not consistent with an isotope effect controlled by the tunnel effect but they are also too large to result from a pure semiclassical effect. The best interpretation would seem again that the shape of the dependence on solvent composition is controlled by the semiclassical isotope effect but with a relatively constant tunnel correction of ca.1.5-2.0 % superimposed. For both Ia and Ib then an interpretation roughly consistent with the calculated curves of fig. 3 is possible. The model used for fig. 3 does seem to underestimate the tunnel effect somewhat and to give it too sharp a maximum at 0.5 extent of hydrogen transfer. A tunnel effect which is nearly constant over the range 0.2-0.6 extent of hydrogen transfer provides better fits to our results. J. BANGER A. FAFFE A.-C. LIN AND w. H. SAUNDERS JR. 119 The limited temperature-dependence data in table 1 also support an important tunnel effect contribution. The observed isotope effects at 80°C in 90 and 100% water are significantly smaller (>95 % probability by the t-test) than those at 60°C though the effects at 60 and 80°C in 60% water are not significantly different.The semi- classical model predicts a somewhat inverse temperature dependence over the 0.2-0.6 extent-of-hydrogen-transfer range while the tunnel correction is predicted to have a normal temperature dependence. 0 20 40 60 80 100 mole % water FIG.4.-Observed k~/k~ values for the reactions of 2-phenylethyldimethylsulphoniumion at 30°C open circles and -trimethylammonium ion at 60°C,closed circles? with hydroxide ion in mixtures of water and dimethyl sulphoxide. One might inquire whether the Bell Sachs and Tranter model l2 can be made to fit our results. Since it gives very small changes in the C-H and 0-H force constants with changes in the attacking base it would predict a nearly constant semiclassical isotope effect.Variations in its predicted tunnel correction depend primarily on changes in the activation energy (strictly speaking the barrier height in whichever direction the process is exothermic but all of the present reactions are exothermic in the forward direction). Activation energies for the sulphonium-salt reaction decrease monotonically as the water content of the medium decreases. While such a change would be consistent with a carbon isotope effect controlled by the tunnel correction it would be incon- sistent with the deuterium isotope effects. The Bell Sachs and Tranter model predicts that both the carbon and deuterium tunnel effects should decrease mono-tonically with decreasing water content.Even if one argues that the overall activation energy is not the appropriate figure for the calculation of the tunnel correction when the proton transfer is part of a more complex reaction coordinate the model would still seem to predict the same dependence on basicity of the attacking base of the tunnel effects for both H/D and 12C/13C. Both are after all of the same origin for the 12C/13Ctunnel correction doubtless reflects the effect of the carbon mass on the tendency of hydrogen to tunnel. In conclusion the Bell Sachs and Tranter model seems to be right in predicting a substantial tunnel correction. It does not seem to be right for the present reactions at least in ascribing to the tunnel effect a dominant role in variations in the isotope effect.This work was supported by the U.S. National Science Foundation. 120 CARBON-13 ISOTOPE EFFECTS L. Melander Isotope Efects on Reaction Rates (Ronald Press New York 1960) pp. 24-32. F. H. Westheimer Chem. Rev. 1961 61 265. R. P. Bell and D. M. Goodall Proc. Roy. SOC.A 1966 294 273. R. P. Bell and B. G. Cox J. Chem. SOC.By1970 194; 1971 783. A. F. Cockerill J. Chem. Soc. By 1967 964. K. C. Brown and W. H. Saunders Jr. unpublished results. E. F. Caldin Chem. Rev. 1969 69 135. R. A. More O’Ferrall J. Chem. Sac. By 1970 785. A. M. Katz and W. H. Saunders Jr. J. Amer. Chem. SOC.,1969 91,4469. W. H. Saunders Jr. Chem. Scripta 1975 8 27. l1 F. G. Bordwell and W. J. Boyle Jr. J. Amer. Chem. Soc. 1971 93 512. l2 R. P. Bell W.H. Sachs and R. L. Tranter Trans. Faraday Soc. 1971 67 1995. l3 R. P. Bell The Proton in Chemistry (Cornell University Press Ithaca New York 2nd edn. 1973) chap. 12. l4 R. P. Bell Chem. SOC.Rev. 1974 3 513. l5 N.-A. Bergman W. H. Saunders Jr. and L. Melander Acta Chem. Scand. 1972 26 1130. l6 W. H. Saunders Jr. unpublished results. l7 M. Wolfsberg and M. J. Stern Pure Appl. Chem. 1964 8 225. l8 J. H. Schachtschneider and R. G. Snyder Spectrochim. Acta 1963 19 117. l9 W. H. Saunders Jr. and S. ASperger J. Amer. Chem. Soc. 1957 79 1612. 2o W. H. Saunders Jr. D. G. Bushman and A. F. Cockerill J. Amer. Chem. Soc. 1968,90,1775. ’’ L. J. Steffa and E. R. Thornton J. Amer. Chem. Soc. 1967 89 6149. 22 W. H. Saunders Jr. and T. A. Ashe J. Ainer. Chem. SOC.,1969 91 4473.23 S. Siggia Quantitative Organic Analysis (John Wiley New York 1963) pp. 457-464. 24 W. H. Saunders Jr. chap. V in E. S. Lewis (ed.) Investigation of Rates arid Mechanisms of Reactions Part I (Wiley-Interscience New York 3rd edn. 1974). 25 G. S. Hammond J. Amer. Chem. Soc. 1955 77 334. l6 C. G. Swain and E. R. Thornton J. Amer. Chcm. SOC.,1962 84 817. 27 E. R. Thornton J. Ainer. Chem. Soc. 1967 89 2915. l8 E. L. Bauer A Statistical Manual for Chemists (Academic Press New York 2nd edn. 1971). l9 i = (-I)+.

ISSN:0301-5696    

DOI:10.1039/FS9751000113

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Structural and Solvent Influences on Tunnelling in Reactions of 4-Nitrophenylnitromethanewith Nitrogen Bases in Aprotic Solvents BY E. F. CALDIN* AND C. J. WILSON University Chemical Laboratory University of Kent Canterbury Kent Received 2nd June 1975 Exceptionally large kinetic isotope effects are observed in the proton-transfer reaction of the carbon acid 4-nitrophenylnitromethane with the strong nitrogen base tetramethylguanidine (TMG) in solvents of low polarity such as toluene ; values of W/kDup to -50 and of AD/AH up to -100 have been reported. The volume of activation has been determined in several of these solvents and shows only small variations. In solvents of higher polarity such as THF dichloromethane or acetonitrile the isotope effects are smaller.With tertiary amines in place of TMG as base the isotope effects show a similar solvent-dependence but are all smaller. These effects are interpreted in terms of tunnelling. The factors influencing the tunnelling correction are discussed (i) the dimensions of the barrier to proton-transfer (ii) possible changes of configuration in the reactant molecules and (iii) solvent motions coupled to the proton-transfer. The data on other reactions exhibiting tunnelling are considered in the light of this discussion. A study is being made 1-4of the influence of variations of the base and the solvent on the kinetics of the proton-transfer reaction of the carbon acid 4-nitrophenyl- nitromethane (RH = NO2.C6H4.CH2NO2)with uncharged nitrogen bases (€3) in aprotic solvents (eqn (1)).The product is deep yellow and convenient for spectrophotometric study. The reaction is shown by p.m.r. to result in a proton-transfer and spectrophotometric determinations of the equilibrium constant are in accord with the scheme (1) in which the product is an ion-pair. The kinetic behaviour also agrees with scheme (1). Rates are conveniently measured by stopped-flow methods '; a wide range of temper-atures can be covered because the activation enthalpies are relatively low. A range of bases can be studied in the less polar solvents 2p especially in chlorobenzene or anisole for which the equilibrium constants are relatively high. Isotope effects are readily measurable.'. 2* 4* Volumes of activation can be determined in some solvents (where the equilibrium constant is ~uitable),~ by measurements of rates at high pressures with a laser temperature-jump apparatus.The rate constants at 25" in a series of solvents increase with solvent polarity; the plots of log kHagainst the reciprocal dielectric constant are somewhat scattered suggesting that specific inter- actions are involved as well as electrostatic forces. The deuterium isotopic effects that have been observed are exceptionally large ; the highest values of kH/kDare about 50 and those of AD/AHare of the order of 100. Such values can only be interpreted in terms of tunnelling. The evidence for tunnel-ling in transfers of H+ H and H-has been reviewed from time to time 6-10 ; a 121 TUNNELLING ; STRUCTURAL AND SOLVENT EFFECTS considerable number of reactions are now known lo which show deviations from classical theory attributable to tunnelling.(Given that atoms have wave properties such deviations are to be expected and the only question is whether they are large enough to be detectable.) The height and width of the potential-energy barrier for pro ton-transfer can be calculated from the temperature-dependence of the is0 tope effect provided we assume a particular shape for the barrier such as a truncated parabola. The base and the solvent have been systematically varied in order to examine their effects on the barrier dimensions. The largest effects are found when the base is tetramethylguanidine or benzamidine containing the grouping HN-C < rather than a tertiary amine ; and when the solvent is of low polarity.In this paper we review the data obtained so far on these reactions and report some new experimental results on the reactions in anisole. The factors that may influence tunnelling effects are considered. EXPERIMENTAL Kinetic isotope effects have been determined by the methods already described '* 2* for the base N,N-diethylbenzamidine (HN=CPhNEt2). The stopped-flow apparatus was the one used earlier.' This work will be more fully described el~ewhere.~ Calculations of the barrier dimensions were carried out by fitting Bell's equations for an unsymmetrical truncated parabola by computer as before.' The original calculations '9 were carried out on two different assumptions (a) that only the proton moves so that MH = 1 and m~ = 2 a.m.u.; (b) that heavy-atom motions are coupled to the proton-transfer so that m~ > 1 with the futher assumption that MD-MH = 1a.m.u.CALCULATIONS In most of the calculations of barrier dimensions reported in ref. (l),it was assumed that the width of the barrier at the base is the same for transfer of H+ and of Df. Because of the different barrier heights for the two isotopes this assumption is not exact; the more correct assumption is that the curvature of the parabolic barrier is the same. Dr. S. Mateo has recalculated his data on this assumption; we are much indebted to him for allowing us to quote his results. RESULTS The results obtained on varying the solvent with TMG as the base throughout are summarised in table 1 which includes new data on the reaction in anisole.The table has three sections (1) solvent properties ; (2) experimentally-determined kinetic quantities and (3) the barrier dimensions calculated from them. For given values of the rate constants the optimum barrier height can be given to k0.05kcal mol-l and the optimum width to +0.002 A. The results obtained on varying the base in a given solvent are summarised in table 2. The results of the calculations of barrier dimensions assuming equal curvature for H and D due to Dr. S. Mateo are shown in table 3 with the results of the previous calculations for comparison. The barrier dimensions are not greatly changed. The new values of the barrier width 2bHare smaller than the previous values of 2b by less than 0.04k The new values of the barrier heights EHand EDare smaller than the previous values by 0.1-0.5 kcal mol-' ; the difference (ED-EH)is usually close to the previous value and is physically reasonable.In general the results confirm the earlier calculations; the differences in 2bH are small and we shall continue to use this as a measure of the barrier width. E. F. CALDIN AND C. J. WILSON 123 TABLE 1.-REACTION OF 4-NITROPHENYLNTTROMETHANE WITH TETKAMETHYLGUAMDINE IN VARIOUS SOLVENTS. RATES,ARRHENIUS PARAMETERS VOLUMES OF ACTIVATION AND BARRIER DIMENSIONS CALCULATED FOR UNSYMMETRICAL PARABOLIC BARRIER cyclohexene mesitylene toluene dibutyl ether anisole D dielectric constant 2.22 2.28 2.38 3.06 4.33 p dipole moment/D 4.5 0.0 0.38 1.22 1.30 a polarisability/cm3moI-' 27 40 30 40 33 103p(n2Va)-' /D mol CM-~ 2.4 0 1.5 3.7 5.2 ETlkcal mol-' 33.0" 33.Y 33.9 35.0 37.2 lO-'k? at 25"C/dm3mol-1s-1 164.5k4 166f5 229 k1 305k8 60723 krf/kD,at 25°C 32.6k1.4 31+1 45k2 41f2 35.7+1 AHJH/kcal mol-' 4.0 &1.4 3.72 &0.03 3.62 k0.05 3.63+0.08 3.2+02 AHlD-AHJHlkcal mol-' 5.4 k0.25 4.7f.0.2 4.3 f0.3 4.2k0.2 --ASJH/cal K-' mol-' 30.4 +0.5 31.3 &0.1 31.0 k0.2 30.4 40.3 30.6 42 log, AD,IA 2.4f.0.2 1.94+0.06 1.5k0.2 1.43kO.13 -3 Vt at 30"C/cm3mol-' -13.2& 1.0 17.8 1.2 -16.34 0.5 -A V" at 30"C/cm3mol-1 -15.9+ 1.7 28.2 &4.6 -29.3 5 1.7 EHlkcal mol-I 11.05 9.8 8.6 8.4 -2b/A 0.820 0.796 0.788 0.786 -267A 0.820 0.796 0.788 0.786 -ni a.m.u.1.oo 1.oo 1.oo 1.oo -chlorobenzene THF dichloromethane acetonitrile D dielectric constant 5.62 7.39 9.08 37.5 p dipole moment/D 1.56 1.7 1.55 3.37 cx polarisability/cin3mol-1 31 18 13 11 103p(n2V")-'/D mol ~m-~ 6.6 12 14 37 ET/kcal mol-1 37.5 37.4 41.1 46.0 10-'ky at 25"C/dm3mol-1s-' 708k20 482k8 52453 589f6 kH,/kD,at 25°C 50+8 12.750.3 11.4f0.2 11.8k0.3 AH!H/kcal mol-1 3.57kO.16 3.44k0.05 3.4940.5 4.3720.04 AHiD-AHJH/kcal mol-I 3.7k0.4 1.8k0.2 1.9f0.2 1.46k0.2 -ASiH/cal K-1 mol-I 28.940.5 30.1 50.2 29.8k0.2 26.640.1 10210 &/AT 1.0k0.3 0.2k0.2 O.35kO.1 O.O+O.1 -A Vt at 3O0C/cm3mol-' 13.04 1.4 -4 V" at 30"C/cm3m01-' 21.9rf 3.5 -EH/kcal mol-' 8.55 4.55 4.85 5.85 2b/A 0.794 0.900 0.954 0.960 2b7A 0.794 0.782 0.786 0.794 m&/a .m.u. 1.oo 1.17 1.24 1.27 E =barrier height ;26 =width of barrier at base ;D =dielectric constant at 25°C ;p =dipole moment ;n =refractive index ;0 =molar polarisability ;ET =empirical solvent-polarity para- meter (some values estimated*).1 cal =4.184 J. DISCUSSION FACTORS INFLUENCING TUNNELLING EFFECTS IN PROTON-TRANSFER REACTIONS In previous discussions 7* of the factors that might favour tunnelling in reactions of carbon acids and the like no simple correlations have been found with the nature of the atoms other than carbon concerned in the reaction (0,F N) nor with the charges on the atoms nor with the electronic properties of the neighbouring groups. It has however been suspected for some time that the effects are influenced by bulky substituent groups l19 l2 and by the solvent.The results summarised above permit TUNNELLING ; STRUCTURAL AND SOLVENT EFFECTS 2.-&UILIBRIUM RATE AND ARRHENIUS TABLE PARAMETERS FOR REACTION OF LGNITRO-PHENYLNITROMETHANE WITH VARIOUS BASES IN TOLUENE ANISOLE AND ACETONITRILE N,N-diethyl-TOLUENE log KHat 25"C/dm3mol-1 TMG t 2.25 benzamidine* quinuclidine 1. 1.61 NEt3 §0.60 NBu~ 6 0.04 -AH"H/kcal mol-1 -AS"H/cal K-I mol-1 10.6f 1.7 25+ 4 8.2f0.6 20+ 2 10.2f0.5 32f1 7.8f0.2 26f1 kr at 25"C/dm3mol-1s-1 2290f 9 2200f 15 132f2 43.okO.4 k"fk; at 25°C 45+ 2 15.6f0.6 11.0k0.7 14+1 AH?"-AH:H/kcalAH,"H/kcal mol-I mo1-l 4.3 +0.3 3.62f0.005 4.5450.09 2.8k0.6 3.51f0.07 5.61fO.14 2.2f0.6 1.7f0.5 -ASFH/cal K-lmol-l 31.Of 0.2 28.0f0.3 37.2f0.2 32.1f0.5 log o(AD,/A';) &/A? 1.50f0.2 32f 14 0.86f0.12 752 0.6k0.45 4f 2 0.04f0.4 1.1 ,QH at 25'C 28f2 4.110.1 3.350.1 3.6k0.1 EH/kcal mol-1 2b/W 2b'/8 m;I/a.m.u.8.6k0.15 0.788 0.788 1.oo 0.96 0.79 1.32 %Oaf 0.05 0.97 0.79 1.30 5.3ofO.05 0.97 0.78 1.35 6.65f0.05 ANISOLE logKH at 25"C/~lm~mol-~ 3.2220.07 -AH"H/kcal mol-' 14.5f 0.8 -AS"H/cal K-lmol-' 38.5f 2.5 1.34f0.05 9.7f0.5 31.3f2.5 2.15f0.02 9.7k0.5 23.0k2 1.05f0.03 10.1k0.9 29.0f3 k,H at 25"C/dm3mol-1s-1 AH,fH/kcal mol-' -AS:H/cal K-lmol-' kp/ky at 25°C 6070k 30 3.2f0.2 35.7+ 1 30.6f2 283f 7 32+ 3 4.2f0.4 33.6f 1 7670f 80 18.8+0.6** 3.25k0.3 30+ 1 620 22.52 1** 3.2f0.4 35-1-1.5 ACETONITRILE log KHat 25"C/dm3mol-1 -AH"H/kcal mot1 -AiS"H/cal K-lmol-l 3.5 4.3 & 0.8 -2f3 2.14 ky at 25"C/dm3mol-1s-1 5890k 60 kTlkp at 25°C 11.8f0.3 AH,"D-AH:H/kcal moP 1.46f0.2 -ASzH/cal K-l mol-' 26.6f0.1 AHf,H/kcal mol-' 4.37 f0.04 1% (AD,IA?) O.O+O.l APIA? l.Ok0.2 12.2k0.9 7.3f0.3 6.7f0.1 6.0f0.1 l.Of0.2 O.Sf0.2 22.150.4 25.2f0.3 1.8f0.3 1.7f0.3 3.12fO.242.17fOO7 0.2~f0.1 0.2~ 0.1 QHat 25°C 2-56 1.67 1.42 EH/kcal mol-l 2b/A 2b'lA 5.8 0.96 0.79 7.7 1.23 0.97 6.85 1.38 0.99 rn&a.m.u.1.27 1.39 1.52 2b"/ArnL/a.m.u. 0.79 1.27 0.79 1.82 0.79 2.05 * This work. 'f Ref. (1). $ Ref. (2). 0 Ref. (5). ** Ref. (4). Symbols and units as in table 1. E. F. CALDIN AND C. J. WILSON a tentative discussion of the influences of the following factors (i) the barrier width and height and the molecular properties that affect them (bonding solvation and steric factors) ; (ii) configurational changes involved in the formation of the transition state and (iii) solvation changes in which the motions of solvent molecules are coupled to the transfer of the proton.(i) THE BARRIER DIMENSIONS The barrier width is the distance that the proton moves in the change from +/ /ICH . .NL to lC ..HN-.\ (It is assumed that the C-N distance does not \/ change appreciably in the very short time concerned.) It must therefore depend (a) on the lengths of the C-H and H-N bonds which probably vary by only a few pm in our systems and (b) on the C-N distance at which proton-transfer occurs which is probably comparable with the length of a hydrogen bond and may be expected to vary appreciably from one system to another depending on the electron- distribution around the nitrogen atom in the base l4 Unfortunately the CH ..N distance has been determined for only one hydrogen-bonded system (HCN),l3.so we cannot argue directly from experimental data. On theoretical grounds however we should expect this distance to be shorter for TMG in which the orbitals of the N atom are sp2-hybridised than for amine bases where the hybridisation is sp3. The overlap integral for the interaction CH ..N will be smaller for sp2 than for sp3 hybridisation unless the distance is less both because the electron density at a given distance from the N atom is smaller and because the orbital is of higher energy; moreover since only two groups are attached to the sp2-hybridised N atom the steric hindrance to formation of a hydrogen bond will be less. The carbon atom will therefore approach more closely to sp2-hybridised N than to sp3-liybridised N before proton-transfer begins to occur.The CH. .N distance will therefore be shorter and the barrier-width smaller for TMG than for amine bases. If the barrier-width depends largely on the type of hybridisa- tion we can understand why the values of 2b calculated from the isotope effects in toluene are about the same for the reactions of all the amine bases regardless of steric considerations and why the value for TMG is appreciably smaller. Steric crowding of the reaction site does not appear to be a major influence on the barrier width which shows only minor variations in the series quinuclidine NEt, NBu, in which the reaction site is progressively more obstructed by the carbon chains.The barrier height will be affected by the changes of bonding by steric factors and by solvation. For TMG and benzamidines the sp2 hybridisation will lead to a shorter C-N distance than that for amines and the additional repulsion energy will increase the barrier height ; we can thus understand the trend in EHfor the reactions in toluene (table 2). For the three amine bases in toluene the variations in EHmay be attributed to steric factors. The barrier height is greater for NBu than for NEt ; this may be attributed to the longer carbon chains in NBu, which will lead to greater repulsion and may also lead to exclusion of solvent molecules from the reaction site so reducing any lowering of the barrier by solvent reorganisation.l0?l2 The still greater barrier height for quinuclidine might arise from a different effect of steric bulk the cage structure of the quinuclidine molecule will prevent the approach of solvent molecules to the N atom in the reaction complex k..H ..N-,/\ and so / could reduce the increase in solvation compared with other amine bases. The less negative entropy of activation is in accord with this suggestion. TUNNELLING ; STRUCTURAL AND SOLVENT EFFECTS Solvent effects on the barrier height (table I) are attributed to solvation partly electrostatic and partly specific (see below).6* lo*12* l6-l8.l9 The effect of asymmetry of the barrier (measured by AH') would be interesting to investigate but at present the data are insufficient. (ii) CHANGE OF CONFIGURATION In the preceding section we have considered only the values of 2b calculated for the various bases by method (a),on the assumption that only the proton moves during the reaction so that the effective mass is that of the proton nzH = 1 a.m.u.If motions of any other nuclei occur and are coupled to the proton-transfer they will contribute to the effective mass along the reaction coordinate and so reduce the tunnelling correction. There may well be changes of configuration of the reacting molecules as a result of the proton-transfer ; for instance the bond angles in NEt and NBu will change slightly when a proton is added.20 Such changes may be coupled with the proton-transfer (though they could alternatively precede the proton- transfer,2' as solvent molecules are thought to do in outer-sphere electron-transfer reactions 22 and in some proton-transfer reactions in water 23).Coupling should in principle be detectable by finding values of the effective mass mfr,by calculations based on method (b)gin which mfr is optimised as well as the barrier width. These I 30 35 40 45 ET/kCal mol-' FIG.1.-Plots of barrier width against the empirical solvent polarity parameter ET for the reaction of 4-nitrophenylnitromethanewith TMG in various solvents. Open circles 2b calculated by method (a) ; filled circles 26' calculated by method (b) ; squares methods (a)and (6); see text. Solvents (2) cyclohexene (3) mesitylene (4) toluene (5) di-n-butyl ether (6) chlorobenzene (7) THF (8) dichloromethane (10) acetonitrile.calculations (table 2) give new values of the barrier width (2b')and of the effective mass (mk) which are just as compatible with the experimental results as are the values of 2b calculated by method (a). For TMG there is no change in the barrier width (2b' = 2b) or the effective mass (mi= 1.00 a.m.u.) ;but for the three E. F. CALDIN AND C. J. WILSON amine bases 2b' is smaller than 2b and nearly equal to the value for TMG while the effective masses mfI are 1.30-1.35 a.m.u. The values of 2b and 2b' are shown in fig. 2. "I e 1' I I I I 2 3 log K FIG.2,-Plots of barrier width against log K for the reaction of 4-nitrophenylnitromethane with various bases in toluene. Open circles 26 calculated by method (a); filled circles 26' calculated by method (b); square methods (a) and (b); see text.Bases (1) TGM (2) quinuclidine (3) triethylamine (5) tri-n-butylamine. TABLE 3.-COMPARISON OF CALCULATED BARRIER DIMENSIONS ASSUMING THAT BARRIERS FOR H+ TRANSFER AND D+ TRANSFER HAVE (a) SAME WIDTH AT BASE (b) SAME CURVATURE FOR REACTION OF 4-NPNM WITH TMG IN VARIOUS SOLVENTS cyclohexane mesi tylene toluene di-n-butly ether barrier parameters (a) (b) (a) (b) (4 (b) (4 (b) CH/kcal mol-l A-" 209 203 172.5 170 208 198 190 182 CD/kcal mol-I A-" 212 210 188 185.5 EH/kcal rno1-l 11.05 10.90 9.80 9.60 8.60 8.45 8.40 9.10 EDlkcal mol-l 11.30 11.15 10.30 9.75 9.40 9.55 8.10 0.818 0.800 0.746 0.744 0.820 0.796 0.788 0.780 0.826 0.814 0.774 0.776 chlorobenzene THF dichloromethane acetonitrile barrier parameters (4 (b) (4 (b) (4 (b) (a) (b) CH/kcal mol-' A-2 165 85.4 76 68 196 86 78.3 70 CD/kcal mol-I A-2 180.5 97.6 86 78 EH/kcal mol-1 8.55 8.45 4.55 4.30 4.85 4.50 5.85 5.55 ED/kcal mol-1 9.75 9.40 5.70 5.40 6.00 5.60 7.00 6.65 2bH/A 0.726 0.882 0.920 0.928 0.794 0.900 0.954 0.960 2bD/A 0.752 0.938 0.978 0.998 CH and CD are curvatures of the parabolas; EHand EDare the barrier heights for the forward reactions ; 2bH and 2bD are the barrier widths.128 TUNNELLING ; STRUCTURAL AND SOLVENT EFFECTS If only NEt3 and NBu were concerned it would be tempting to adopt the results of method (b) and explain the high values of mfr as due to configurational changes. It is difficult however to explain the results for quinuclidine and TMG on this basis.The effective mass m;I calculated for quinuclidine by method (b) is about the same as for NEt and NBu, whereas it should be smaller if due to configurational changes which must be hindered in quinuclidine. For TMG the calculated effective mass mfr (1.00 a.rn.11.) would imply that there is no coupled configurational change although protonation of TMG leads to a change of hybridisation which must alter bond lengths and angles and so affect mfI if configurational change is relevant. These anomalies in the results of method (b)make us incline to the view that the true barrier widths are those given by method (a). with m;I = 1.00 a.m.u. so that there is a real difference between the barrier widths for TMG and for amine bases; and that there is no clear evidence from the present results for effects of coupled configura- tional changes.(iii) SOLVENT MOTIONS COUPLED TO PROTON-TRANSFER The formation of a dipolar transition state in a solvent of low dielectric constant would be expected to lead to reorientation of solvent molecules. Rotation of solvent inolecules might therefore be coupled to the proton-transfer.l* l2 This would increase the effective mass. The results for TMG in table 1 section 3 and fig. 1 show a marked difference between the less polar and the more polar solvents. (Effects with other bases (table 2) are smaller but similar.) For the reaction with TMG in the less polar solvents (dielectric constant < 6) the barrier widths 2b calculated by method (a) i.e.assuming m = 1.00 a.m.u. are nearly constant (0.80+0.02 A) ; the same values are obtained for 2b’ by method (b) and the optimum value of the effective mass remains 1.00 a.m.u. It appears there- fore that in these solvents there is no coupling of solvent motions with the proton- transfer. This conclusion is also in accordance with the values of the volumes of activation (AV*) for the reaction of TMG determined from rate measurements at high pre~sure.~ In mesitylene toluene anisole and chlorobenzene AV* is approximately constant (table 1) ; all the values lie within the range -15k3 cm3 mol-I. A value of this order (ca.-12 cm3 mol-I) is obtained by a rough calculation from a simple model for the decrease in volume of the reacting pair of molecules when the CH .. N distance decreases from the van der Waals distance to the hydrogen-bonded distance omitting all consideration of solvation changes. The values vary less than the overall volumes of reaction AVO in the several solvents which range from -16 to -29 cm3 mol-’ ; and they are smaller than the values of AV* for typical Menschutltin reactions,24* 25 in which charge-separation associated with transfer of heavier nuclei occurs (-20 to -50 cm3 mol-I). This evidence is compatible with the view that in these solvents the proton-transfer is neither accompanied nor preceded by any considerable rotation of solvent molecule?. It does not support the view that solvent molecules move first and the proton-transfer occurs when the solvent environment is suitable for the product as in outer-sphere electron-transfer reactions in water.For the reaction of TMG in the more polar solvents (THF dichloromethane acetonitrile) calculations by method (a) lead to values of the barrier width 2b markedly larger than for the less polar (0.90-0.96 A). Since there is no obvious physical reason for so large a difference (though the effect of dielectric constant on the initial hydrogen-bond distance CH . .N may account for part of it) calculations were carried out by method (6). These gave optimum values of the barrier width in close agreement with those found for the less polar solvents (26’ = 0.78k0.01 A) E. F. CALDIN AND C. J. WILSON and the corresponding values of the effective mass are mk = 1.17 to 1.27 a.m.u.(table 1 section 3b). Such values can be understood in terms of coupling of the reorientation of solvent molecules with the proton-transfer. These solvents are the ones in which the field due to a dipolar transition state on adjacent molecules will be largest. This field according to a simple model due to Dr. N. J. Bridge (personal communication) will be proportional to p/Von2,where p is the dipole moment of the molecules of solvent Yoits molar volume and n its refractive index; the values of this function (table 1 section 1) are much greater for the three solvents mentioned than for the less polar ones. The numerical values found for rnI; are in reasonable agreement with those expected from a simplified electrostatic model of the coupled motion due to Prof.R. P. It is possible that another factor influencing the behaviour of the three more polar solvents is that each of them has a relatively low moment of inertia about at least one axis while the less polar molecules have not. Experimental difficulties have so far prevented us from extending the range of solvents to investigate this point. It was also found impossible to determine volumes of activation in the more polar solvents because the product was not stable for the relatively long periods required. The picture that emerges for the reaction with TMG in the more polar solvents is that when the proton-transfer occurs the field due to charge separation in the transition state exerts a torque on the polar solvent molecules so that they begin to rotate; their motion is therefore coupled to the transfer of the proton and this coupling leads to an increase in the effective mass which is reflected in a smaller tunnelling correction and hence in smaller isotope effects.For the reaction in the less polar solvents the corollary of this view is that for some reason the fleld due to the transition state does not bring about rotation of solvent molecules although some interaction is indicated by the variations of the barrier height EH. The reason may be that in these solvents the effect of the field is only to bring about electron polarisation in the solvent molecules rather than rotation. These solvents have relatively high molar polarisabilities and relatively low values of the quantity ,u/n2/Vorepresenting the field (table l) compared with the more polar solvents.Electron polarisation is set up in -10-I' s and will be coupled with the proton-transfer but will exert a negligible torque and will leave the effective mass unaltered. REQUIREMENTS FOR LARGE TUNNELLING EFFECTS IN REACTIONS OF CARBON ACIDS The main factors that produce large tunnelling effects in reactions such as that of 4-NPNMwith TMG in certain solvents thus appear to be (1) the unusually narrow barrier attributable to the sp2 hybridisation of the orbital of the basic N atom; (2) the barrier height which may be increased by steric bulk in the base; (3) the low polarity of the solvent which appears not to undergo rotations coupled to the proton- transfer and (4) the apparent lack of a configurational change coupled to the proton- transfer.It is of interest to enquire whether these factors also characterise the other reactions that exhibit anomalous isotope effects attributable to lo and what other factors may be involved. (1) Evidence on the width of the barrier in a given reaction is not generally available ; indeed part of the interest attaching to tunnelling calculations is that they provide such evidence. But if it is accepted that the barrier is narrower for sp2-hybridisation of the nitrogen orbitals than for sp3-hybridisation then it may be significant that of the 18 proton-transfer reactions (excluding that of 4-NPNM with S 10-5 130 TUNNELLING ; STRUCTURAL AND SOLVENT EFFECTS TMG) listed in Bell's tables of reactions lo showing anomalous hydrogen isotopic effects on Arrhenius parameters four are reactions of pyridine bases.(2) As regards the barrier height it may be significant that five are reactions of hindered bases such as collidine or t-butoxide. (3) Reactions in which the solvent reorganisation is small or unusually easy will be favourable cases for large tunnelling effects. (a) It is noteworthy that 10 of the proton-transfer reactions listed do not involve a change of charge on forming the transition state so the solvation change will be less than in an ionogenic reaction. (b) The solvent is aqueous for 16 of the reactions including 6 of the 8 ionogenic reactions other than 4-NPNM +TMG. Water is unique among common solvents in the small moment of inertia of its molecules ; small motions in water appear to be almost ~nhindered,~~ and solvent reorganisation will have relatively little effect on mA.(c) The effect of steric bulk in the base molecule may be to exclude solvent from the reaction site (cf. above) and so reduce effects due to solvent reorganisation. Hydrogen-atom transfers are particularly favourable from this point of view since they involve no charges and can be carried out in non-polar solvents. An example is the reaction of oxygen (0,)with dihydrophenanthrene in a hydrocarbon which has kH/kD= 64 at -10°C corresponding to about 30 at 25"C and has been the subject of a very full theoretical investigation. The reactions of sub-stituted phenols with radicals in vinyl acetate give kH/kDup to 19 at 50°C.18* 29 Eighteen such reactions are listed by Bell.O (4) Evidence on configurational changes might be sought by comparing rigid with non-rigid systems (though the results for the reactions of 4-NPNM with bases have so far been negative). It is of interest that the proton-transfer reactions of 2-carbethoxycyclopentanone 30 (the first established case of a large tunnelling correction) and the H-atom transfer reaction of dihydrophenanthrene 28 involve ring systems. CONCLUSION The general results of this investigation of the reactions of 4-NPNM with nitrogen bases so far may be summarised as follows. (1) The isotope effects vary widely with base and solvent. The largest are explicable only in terms of tunnelling. (2) When the base is varied much larger effects are found with tetramethyl- guanidine or diethylbenzamidine than with amine bases.The larger effects can be attributed to a smaller barrier width due to the more compact electron-distribution in the sp2-hybridised orbital. Among the amine bases there is little difference in the calculated barrier dimensions between quinuclidine and NEt or NBu3 ; this suggests that configurational changes accompanying proton-transfer do not alter the effective mass appreciably. The main factor influencing the effect of a base thus appears to be the electron-distribution around the nitrogen atom. (3) When the solvent is varied the largest effects are found in solvents of low polarity (B< 6). The explanation proposed is that in the more polar solvents the solvent molecules interact with the field set up by the separation of charges in the transition state and begin to rotate so that their motion is coupled to that of the proton ; the effective mass is thereby increased and the tunnel effects decreased.In the less polar solvents which are also the more polarisable we may suppose that only electronic polarisation occurs so that the effective mass is nearly that of the proton and the tunnelling effects are not appreciably reduced. The main factor influencing the effects of the solvent appears to be its polarity. The size of the molecule could be significant. E. F. CALDIN AND C. J. WILSON 131 (4) When we consider the data on other reactions which involve transfer of Hf PI or H-to or from carbon and have large isotope effects attributable to tunnelling it appears that the expected effects of the solvent and of bulky substituents can be seen but that effects of configurational change are not clearly discernible.We acknowledge helpful discussions with Prof. R. P. Bell Dr. N. J. Bridge Prof. C. D. Hubbard and Prof. F. Ann Walker and an S.R.C. research fellowship to C. J. W. E. F. Caldin and S. Mateo J.C.S. Faraday Z 1975 71 1876. E. F. Caldin and S. Mateo J.C.S. Faraday Z 1976,72 112. C. D. Hubbard C. J. Wilson and E. F. Caldin J. Amer. Chem. SOC.,1976 98. E. F. Caldin D. M. Parbhoo C. J. Wilson and F. A. Walker J.C.S. Faraday I 1976,72. E. F. Caldin A. Jarczewski and K. T. Leffek Trans Faraday SOC. 1971 67 110. R. P. Bell The Proton in Chemistry (Chapman and Hall London 1973) chap.12. E. F. Caldin and M. Kasparian Disc. Faraday SOC.,1965 39,25. E. F. Caldin Chem. Rev. 1969 69 135. 9 E. F. Caldin in Reaction Transition States ed. J. E. Dubois (Gordon and Breach London 1972) p. 247. lo R. P. Bell Chem. SOC.Rev. 1974 3 513. l1 E. S. Lewis and L. H. Funderburk J. Amer. Chem. SOC.,1967,89,2322. E. S. Lewis in Proton-transfer Reactions ed. E. F. Caldin and V. Gold (Chapman and Hall London 1975) chap. 10 l3 G. C. Pimentel and A. L. McClellan The Hydrogen Bond (Freeman San Francisco 1960). l4 S. N. Vinogradov and R. H. Linnell Hydrogen Bonding (van Nostrand New York 1971) chap. 7. l5 R. D. Green Hydrogen Bonding by C-H Groups (Macmillan London 1974). F. H. Westheimer Chem. Rev. 1961 61 265.l7 J. R. Keeffe and N. H. Munderloh Chem. Comm. 1974 17. l* M. Simonyi I. Fitos J. Kardos I. Lukovits and J. PospiSil Chem. Comm. 1975 252. l9 R. A. More O’Ferrall in Proton-transfer Reactions ed. E. F. Caldin and V. Gold (Chapman and Hall London 1975) chap. 8. 2o H. C. Brown H. Bartholomay and M. D. Taylor J. Amer. Chem. SOC.,1944 66,435. 22 J. L. Kurz and L. C. Kurz J. Amer. Chem. SOC.,1972,94,4451. 22 R. A. Marcus Ann. Rev.Phys. Chem. 1964,15,155; J. Phys. Chem. 1968,72,891. 23 (a) W. J. Albery A. N. Campbell-Crawford and J. S. Curran J.C.S. Perkin ZZ 1972 2206 ; (6) M. M. Kreevoy and D. E. Konasewicli Adv. Chem. Phys. 1971 21,243 ; (c) M. M. Kreevoy and Sea-wha Oh J. Amer. Chem. SOC. 1973,95,4805 ; (d)A. J. Kresge Chem. SOC.Rev. 1973,2,484 ; (e) M.H. Davies J. R. Keeffe and B. H. Robinson Ann. Rep. Chem. SOC. A 1973 123 ; (f)W. J. Albery in Proton-transfer Reactions ed. E. F. Caldin and V. Gold (Chapman and Hall London 1975) chap. 9. z4 H. Hartmann H. D. Brauer H. Kelm and G. Rinck Z.phys. Chem. (Frankfurt) 1968,61,53. 25 H. Heydtmann A. P. Schmidt and H. Hartmann Ber. Bunsengesphys. Chem. 1966 20,444. 26 R. P. Bell personal communication. 27 (a) D. W. G. Smith and J. G. Powles Mol. Phys. 1966 10 451 ; (b)H. G. Hertz Aizgew. Chem. (Int. Ed.) 1970 9 124; (c) D. Eisenberg and W. Kauzmann The Structure and Properties of Water (Clarendon Press Oxford 1969). p. 214. 28 A. Bromberg K. A. Muszkat E. Fischer and F. S. Klein J.C.S.Perkin ZI 1972 588 and earlier papers. 29 M. Simonyi and F. Tudos Adv.Phys. Org. Chem. 1970 9 127 ; 3O K.P. Bell J. A. Fendley and J. R. Hulett Proc. Roy. SOC.A 1956 235 453. S10-5*

ISSN:0301-5696    

DOI:10.1039/FS9751000121

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

GENERAL DISCUSSION Dr. R. A. More O'Ferrall (University College Dublin) said :The initial formation of a tetrahedral carbanion in the ionisation of nitromethane would seem to imply that the main activating effect of the nitro group is inductive or electrostatic. Why then does hydrogen exchange occur so much more readily than in a tetramethyl-+ ammonium ion?l A (CH&N substituent is unable to conjugate but since it is a positive monopole its inductive effect should be greater than that of the dipolar NO2 group. Prof. F. G. Bordwell (Northwestern University) said The rate of exchange for the tetramethylammonium ion may be much slower than the actual rate of carbanion formation because of internal return. The large kH/kDratios observed for nitro- alkanes indicate that internal return is not a problem in this instance.Furthermore we believe that the effect of the nitro group on the formation of the "essentially pyramidal "nitro carbanion is not purely electrostatic in nature but instead involves conjugation. Even a pyramidal anion should overlap to some degree with the n-system of the nitro group; such overlap should be appreciable in intermediate 2 if as we have suggested some rehydridization and solvent reorganization has occurred in its formation. Dr. W. J. Albery (Oxford University) said First I would like to ernphasise the similarity between the diazo systems discussed earlier and the nitro system i B-+ H -C -NO2$B---H 1 \ 0- W. von E. Doering and A. K. Hoffman J. Amer. Chem. SOC.,1955,77 521.*W. J. Albery A. N. Campbell-Crawford and J. S.Curran J.C.S. Perkin 11 1972 2206. 132 GENERAL DISCUSSION Secondly while we may write a tetrahedral intermediate such as (2) we do not intend to imply that the system moves all the way to (2). If one believes the Marcus theory it may be shown that the most favourable path for reaction is when the thermodynamic driving force for the proton transfer (AG;’ in Marcus terminology and AGO’ in Kreevoy’s notation ’) is equal to zero. Thus in both cases the amount of lone pair development in (2) should be such that the pK of the organic substrate roughly corresponds to the pK of the cataly~t.~ Prof. A. J. Kresge (University of Toronto)said The anomalous Bronsted exponents observed in nitroalkane ionization imply that negative charge builds up on the a-carbon atom in the proton transfer transition state of this reaction and explanation of this unexpected behaviour therefore reduces essentially to justifying the greater negative charge density at this atom in this transition state than in the stable nitronate- ion reaction product.Bordwell’s justification postulates a new reaction intermediate a second anionic species from which the proton has been transferred completely but in which the electron pair left behind has not yet moved to a more stable location. This explanation would seem to have electrons moving more slowly than protons. It is possible however to avoid this difficulty by assuming that complete electroe reorganization requires a completely formed carbon-nitrogen double bond ; in the proton transfer transition state this double bond is only partly formed and a portion of the negative charge generated in the substrate therefore must remain localized on its a-carbon atom.4 This explanation provides a mechanism whereby electron delocalization may lag behind proton transfer in a sin& reaction step and thereby removes the need for inventing a new reaction intermediate.Prof. F. G. Bordwell (Northwestern University) said The conversion of 2 to 3 involves not only “ movement of electrons to a more stable location ” but as pointed out in the paper partial rehybridization (movement of atoms) partial solvent reorganization (movement of molecules) and breaking of the strong BH . . . C-H-bond.Each of these factors is believed to contribute to the activation barrier between 2 and 3. Intermediate 2 is stabilized by solvation forces and strong H-bonding. Prof. F. G. Bordwell (Northwestern University) said Bell pointed out in his lecture that compounds which differ appreciably in structure often do not fall on the same Brijnsted plot. Nitroethane and 1,l-dinitroethane are likely to fall in this category. A study in our laboratory (J. E. Bartmess unpublished results) has shown that for a series of substituted nitroethanes of the type G(CH,),CH2N02 p* = 1.2 for equilibrium acidities in 50 % (vlv) MeOH-H20 (p* E 1.0 estimated for H20). Data of Sitzmann Adolph and Kamlet give p* = 3.60 for equilibrium acidities of a comparable series of 1,1-dinitroethanes G(CH2),CH(N02)2in water.This means that the negative charge in the dinitro anions must be concentrated to a much greater degree on carbon (and nitrogen) then is true in the mononitro anions. This is contrary to what might have been expected on the basis of the relative acidities in the two series. The data point to the presence of strong steric inhibition of resonance which renders difficult a straight-forward interpretation of solvent effects. R. A. Marcus paper at this Symposium. A. I. Hassid M. M. Kreevoy and T. Laing paper at this Symposium. W. J. Albery A. N. Campbell-Crawford and J. S. Curran J.C.S. Perkin 11 1972 2206. A. J. Kresge Canad. J. Chem. 1974 52 1897. M. H. Davies J.C.S. Perkin II 1974 1018. M. E. Sitzmann H. G. Adolph and M.J. Kamlet J. Amer. Chem. Soc. 1968,90,2815. GENERAL DISCUSSION Dr. B. G. Cox (University of Stirling) said The reactions of nitroethane and 1,l dinitroethane with acetate (and hydroxide) ions have been quoted as examples of reactions showing a Bronsted coefficient greater than unity. Our results simply show that in accordance with expectations based on the relative solvation of the mono- and dinitroalkane anions the Bronsted coefficient decreases (and becomes “normal ” i.e. 0 < p < 1) as the DMSO content of the solvent increases. With reference to the results quoted by Bordwell an alternative explanation of them may be given again in terms of the different solvation of the anions. Thus the much stronger interactions with protic solvents of the mono nitro alkane anions would be expected to reduce the p* values of the nitroethane series relative to the 1,l dinitroethanes.The p* values should be much closer and perhaps even reversed in dipolar aprotic solvents or in the gas phase. Dr. B. R. Eggins (N. Ireland Polytechnic) said I think it is a pity that Cox has not extended his experiments in DMSO +water mixtures to very high mole fractions of DMSO including very dry DMSO and to very low fractions of DMSO. We have observed some unexpected effects in the oxidative dephosphorylation of hydroquinone phosphates. The relative rates of P-0 and C-0 bond fission during the oxidation of durohydroquinone monophosphate in mixtures of alcohols and water was found to be relatively constant at 39f2 % in methanol ethanol and propan-1-01 up to 0.8 mole fraction of alcohol which is similar to the results found by Kirby and Varvoglis and by Bunton Fendler and Fendler for the hydrolysis of p-nitrophenylphosphate.It is probable that the rate of proton transfer is a key step in determining these relative rates.5 However in solutions of mole fraction of alcohol greater than 0.8 the rate ratios decreased sharply. This decrease was also observed with propan-2-01 and trifluoroethanol. Our suggested explanation for this effect is that it is due to the change in the structure of solvent water as the diamond lattice can be completely broken down when four molecules of another suitable solvent can solvate each water molecule.6 The fast proton transfer which is possible by proton jump transfer in structured water does not occur when the structure is lost.As further evidence for this hypothesis the effect is partially reversed when very dry methanol is used and there as an increase in the percentage P-0 bond fission. This may be due to the restoration of linear structure to methanol which can allow limited proton jump transfer. While not wishing to deride the enormous value that linear free energy relation- ships have provided in correlating kinetic and thermodynamic data and showing up structural effects in reactions I feel that in attempting to explain anomalies in Bronsted plots such as curvature and p values greater than unity it is necessary to consider other factors more fully. By their very nature linear free energy relation- ships are energy difference plots and are intended to eliminate factors which are assumed to be constant throughout a reaction series.However this means that the effects due to changes in such factors are masked. I suggest that in Cox’s work and that of others involving mixed solvents changes of solvent structure should not be ignored. R. P. Bell and R. L. Tranter Proc. Roy. Soc. A 1974 337 517. B. R. Eggins and D. W. Hutchinson unpublished results referred to in V. M. Clark and D. W. Hutchinson,Prog. Org. Chem. 1968,7 100. A. J. Kirby and A. G. Varvoglis J. Amer. Chem. SOC., 1967 89,415. C. A. Bunton E. J. Fendler and J. H. Fendler J. Amer. Chem. Soc. 1967 89 1221. W. W. Butcher and F. H. Westheimer J. Amer. Chem. Soc. 1955 77 2420.F. Franks and D. J. G. Ives Quart. Rev. 1966,20 1. GENERAL DISCUSSION Prof. J. B. Wyne (University of Calgary) said I would like to comment on the papers of both Cox and Gibson and of Saunders and his co-workers. In both of these studies aqueous organic binary solvent mixtures were employed in order to effect variation of the bulk properties of the solvent medium but little or no comment is made regarding the important and sometimes highly complex variations in these systems at the molecular interaction level. Cox and Gibson report that the rates of proton transfer from nitroalkanes change in the opposite direction as the mole fraction of organic cosolvent is increased in water +dimethylsulphoxide (DMSO) compared with water +2,2,2-trifluoroethaiiel (TFE).They rationalise this behaviour in terms of the large difference in the interaction of water with anions of high charge density. In our recent examination of the heats of solution of quaternary ammonium salts in aqueous binary mixtures we have measured the AHs of tetramethylammonium chloride in both water +DMSO and water +TFE mixtures. The results shown in fig. 1 clearly establish that there are not only large differences in the solvent-ion interactions in these two systems but the heats of solution actually change in the opposite direction as observed by Cox and Gibson for the rates of proton transfer. AH of Mc,,NCI in Voii3us Aqueous Orgunic E.lix!ures ot 25°C FIG.1. These authors also note that anomalies in the Bronsted coefficient are character- istically associated with the high water content end of the solvent mixture range.This of course is the region in which the added organic cosolvent has its maximum effect on the unique "structuredness " of water. The AHs data in fig. 1 reflects this structural influence in the 0.0 to 0.1 mole fraction organic cosolvent region and in the case of water -I-TFE mixtures the AHs against mole fraction dependence actually changes sign. Again in the elegant and exhaustive kinetic study of carbon-13 isotope effects on proton transfer Saunders and his co-workers have employed dimethylsulphoxide + GENERAL DISCUSSION water mixtures. Re-examination of the plots in fig. 1 and 4 of their paper indicate that in all cases reasonable curves can be drawn through the experimental points indicating extrema at approximately 0.33 mole fraction dimethylsulphoxide.This is the same composition as has been frequently observed to characterise the extremum behaviour in many other phenomena in DMSO +water mixtures including excess heats and volumes of mixing and activation energies for various hydrolyses in these solvent mixtures. Parker and others have agreed that at this composition the unique structural features of water can be assumed to have been completely destroyed by the organic cosolvent and that a loose 2 1 H20:DMSO complex characterises the solvent system. While it is correct to say that this extremum behaviour in aqueous organic media is not normally found in free-energy ((AG log k etc.) composition) plots due to enthalpy/entropy compensation (Lumry’s Law) plots of the ratio of free energy terms as is the case in fig.1 and 4 in Saunders’ paper could well exhibit such behaviour. In any event the occurrence of extrema in the isotope effect against solvent composition plots at the same mole fraction as has been ascribed to the formation of a particular DMSO+water complex and as is observed with many other properties deserves some attention. It is also known that small additions of DMSO (among many other organic cosolvents) to water has a very marked effect on the so-called “structure ” of water. This is an important point in explaining the observed difference in isotope effect between 100 % and 95 % water which surprised Saunders and his co-workers.Very substantial changes in ionic heats of solution occur between pure water and 0.05 mole fraction DMSO. Accordingly it is not surprising that the kinetic isotope effect for proton transfer presumably via some form of Grotthus chain mechanism should undergo considerable change as the structure of the solvent water is effected by initial additions of organic cosolvent . Dr. B. G. Cox (University of Stirling) said In replying to Hyne and Eggins I would like to emphasise that a major point of our paper is that the effect of solvent variation on the rates of the proton transfer reactions studied is not simply related to (and may be opposite to) the effect on the corresponding equilibria. This shows that solvent effects whether they arise from “structural ” effects or H-bonding interactions etc.do not in general vary monotonically with the extent of proton transfer during the reactions and hence will influence observed p values. We have studied the reactions reported in mixed solvents with the mole fraction of the organic component varying between 0.06 and 0.85 and found no extremain the variation of rates with solvent composition. We feel that in DMSO+H20 water mixtures in particular the very large rate increases observed (>lo3) on transfer from water to the mixtures can most simply be interpreted in terms of the desolvation fo the anion bases (OAc- F- R,CNO;). Prof. W. H. Saunders (Uaiversity of Rochester) said :We assumed in interpreting our results that addition of dimethyl sulphoxide to aqueous hydroxide changes rates and isotope effects mainly by increasing the basicity of both reactant hydroxide ion and hydroxide ion in the transition state.We cannot exclude the possibility that other solvent effects enter in but there is in our opinion no definite evidence for such effects. As Hyne points out free energies need not reflect extrema in enthalpies. Our isotope effects reflect differences in free energies of activation. Since the rate for the light isotopic species changes monotonically with solvent composition it is hard to see why a solvent effect unaccompanied by a change in the extent of proton transfer would not change the rate for the heavy isotopic species in a parallel manner. GENERAL DISCUSSION The H-function for the medium does correlate well with rate which rises mono- tonically with increasing concentration of dimethyl sulphoxide and shows no pecularities in either high-water of 2 1 water dimethyl sulphoxide regions.Turning to several specific points raised by Hyne the minima in the carbon isotope effects (fig. 1) are shallow and difficult to locate but do not appear to occur at the same solvent composition. Neither do the more easily located maxima in kH/kD (fig. 4) so that clear evidence for specific solvent effects in the 2 1 water dimethyl sulphoxide region is lacking. As for the possible relation between changes in water structure and the scatter of carbon isotope effects for the sulphonium salt in 95-100 % water the slowness of the reactions and our failure to control sample sizes closely enough could explain the scatter.Consequently I do not feel justified in yielding to the temptation to attribute the scatter to specific solvent effects. Dr. W. J. Albery (Oxford University) said In presenting his paper Marcus com- pared electron and proton transfers and emphasised that while electrons may tunnel through distances of up to 1 nm a simple proton transfer takes place over a much shorter distance. For this reason he expected that the solvent oscillation which facilitates electron transfer may be less important in proton transfer. However while this argument may be true for simple proton transfers as soon as one has a carbon base with a n system then the movement of charge out of for instance the diazo or nitro group may provide sufficient leverage for solvent oscillations to be significant for these systems.This would explain why these systems have substantial Wand Wp terms. Prof. R. P. Bell (University of Stirling) said It seems to me rather artificial to speak of the effect of the isotopic mass of carbon upon the tunnelling of hydrogen since in the transition state the system can be represented by a single reduced mass moving along the reaction coordinate. It would be of interest to know how much this reduced mass depends upon the extent of H-transfer and also how much it varies as the system moves away from the transition state. The latter variation could be significant when the tunnel correction is considerable. Prof. W. H. Saunders (University of Rochester) said It is of course just a way of looking at a process which cannot really be dissected in this fashion but it is not entirely without mechanistic merit.A large tunnel correction requires a reaction coordinate frequency of fairly large absolute magnitude and fairly large isotopic sensitivity. The dependence of frequency on reduced mass suggests that motion of hydrogen must contribute substantially to the reaction coordinate for the first condition to be satisfied. Furthermore the model calculations on the E2 reaction give a sizable tunnel correction to the P-carbon isotope effect a considerably smaller one to the a-carbon isotope effect and essentially none at all to the leaving-group isotope effect even for 14Nagainst 5N.Thus conditions favourable to a significant tunnel correction to a heavy-atom isotope effect seem to include direct coupling of hydrogen motion to the heavy-atom motion.On the other hand it must be admitted that the sensitivity of the reaction coordinate frequency to isotopic substitution at carbon will be least when the reduced mass is closest to the mass of hydrogen a condition which will hold when the C-H and 0-H stretching force constants are nearly equal. A simple triatomic model yields a negligible tunnel correction under these circumstances. The E2 model does better because the reaction coordinate always involves more heavy-atom motion. The isotopic sensitivity of the reaction-coordinate frequency rises for less symmetrical GENERAL DISCUSSION transition states since the extreme reaction coordinates for the triatomic model can be regarded as approach of 0to CH and retreat of OH from C respectively.Prof. P. Zuman (Clarkson College) said Differences between the dependence of H-function (obtained for dissociation of anilines or indoles) and J-function (obtained for addition of OH-ions to benzaldehydes) on DMSOlwater ratio indicates that H-is not a simple function of OH-hydration. Change of H-function with DMSO concentration should thus not be used as a measure of OH-hydration. Prof. W. H. Saunders (University of Rochester) said E. S. Lewis’ proposal that steric hindrance promotes tunnelling suggests that a larger tunnel effect ought to be observed with triethylamine and tributylamine than with quinuclidine which has the alkyl groups on nitrogen “tied back ”.Yet Caldin’s data in table 2 indicate the tunnel effect is largest with quinuclidine. Please comment on this departure from the expected order. Prof. E. F. Caldin (University of Kent) said The tunnelling correction (QH) at 25°C in toluene for the reactions of 4-nitrophenylnitromethane decreases in the order quinuclidine > tri-n-butylamine > triethylamine. The curvature of the bar- rier decreases in the same order; this is because the barrier height decreases while the width varies little. The higher barrier for tri-n-butylamine compared with triethylamine can be interpreted in terms of steric hindrance either (1) by the original idea that the effect of bulky groups is to increase the repulsive forces between reactant molecules or (2) by E.s. Lewis’ recent suggestion that their effect is to hinder solvation of the transition state thus reducing the effects of coupling of the proton transfer with motions of heavy solvent molecules which would otherwise increase the effective mass and SO reduce the tunnelling correction. To account for the still higher barrier height for quinuclidine in these terms it is necessary to suppose that exclusion of solvent on the side remote from the proton is important ; the cage structure of the quinuclidine molecule makes it impossible for the solvent to approach the nitrogen atom in the reaction complex (C . .H . .N-from this side whereas 3 with the other amines there is much less hindran~e.~ Dr.W. J. Albery (Oxford University) said Accepting Caldin’s analysis I would like to ask how much coupling does mfI = 1.3 represent? Could it be that tunnelling is SO easily quenched by even a little heavy atom motion that the Franck Condon separation of H+ and solvent motion is still a good approximation? Prof. E. F. Caldin (University of Kent) said In reply to Albery one way of quantifying the extent of coupling in proton-transfer represented by an effective mass of 1.3 a.m.u. (rather than 1.0 a.m.u.) is to calculate the effect on the tunnelling correction Q,which is the ratio of the rate constant to that which would be observed with the same potential-energy barrier if there were no tunnelling (i.e. if the proton could be treated as a particle rather than a wave).We use Bell’s equations and simplify by taking only the first term and by assuming AH” = 0. Then Q = +u/(sin 3u) ; u = hv,/kT; v = Et/nb (2m).f. 1 L. H. Funderburk and E. S. Lewis J. Amer. Chern. Soc. 1964 $6,2531. 2 E. s. Lewis in Proton-transfer Reactions ed. E. F. Caldin and V. Gold (Chapman and Hall London 19751 p. 333. 3 E. F. Caldin and S. Mateo J.C.S. Furaduy I 1976,72 112. GENERAL DISCUSSION Let Q refer to rn = 1 a.m.u. and Q’ to m = 1.3 a.m.u. and suppose the barrier dimensions E and 2b are constant. Then taking the value of u as for acetonitrile (4.3) we find Q/Q’ = 1.4. Thus coupling of solvent motions with proton transfer to the extent that increases the effective mass to 1.3 a.m.u. does not reduce Q by a factor comparable with Q,which for several of the solvents in our work is over 20 at 25°C (see ref.(l) of our paper).For toluene for instance QHat 25°C is 28 and would be reduced to 20. Most of the reduction in QH(and therefore in kH/kD) on passing from toluene to acetonitrile for instance (for which QH = 2.6 at 25°C) is due to the decrease of curvature of the barrier (v = 1 420 cm-l for toluene 956 cm-1 for acetonitrile) which in turn is due to the decrease in barrier height (EH= 8.60 for toluene 5.85 kcal n101-~ for acetonit- rile). This decrease in EHis attributable to increased solvation of the transition state whether due to electrostriction or to specific interactions. Prof. R. P. Bell (University of Stirling) said It is worthwhile emphasizing that the calculated barrier dimensions in tables 1 and 2 involve the assumption m,-mH = 1.This is a somewhat arbitrary assumption though it does follow from a highly simplified electrostatic model of the coupling between proton motion and solvent rotation; however I do not think that any refinement of this treatment is likely to lead to essentially different results. Dr. W. J. Albery (Oxford Unirersity) and Prof. PI. M. Kreevoy (Uuiveusity of Minnesota) (communicated) We would like to offer an explanation for the variation in EHwith solvent observed by Caldin and Wilson. In the more polar solvents specific solvent-solute interactions can stabilize the transition states in which there is already some charge separation. This leads to a reduction in EH,the barrier to the actual proton transfer (see table I of Caldin’s paper).In terms of the Marcus theory this corresponds to a reduction in AG’ but at the expense of an increase in W‘. Since only AG# is susceptible to tunnelling the barrier height calculated from the manifestations of tunnelling goes down dramatically even though the reaction A FIG.1.-Effect of changing the solvent from a non polar solvent A to a polar solvent B on thc truncated parabolic energy barrier to proton transfer. On the left are shown the corresponding free energy terms of the Marcus theory ; W*must be small for the non polar solvent A but can be much larger for the polar solvent B. GENERAL DISCUSSION does not become much faster. The situation is shown schematically in fig.1 in which following Caldin's argument we have kept the barrier width constant. This explanation accommodates very nicely the general observation that tunnelling corrections are small in polar solvents even when the large primary isotope effect suggest that the effective mass in the reaction coordinate must be close to the mass of the proton. The flat squat barrier of curve B will only lead to a small tunnelling correction. The pattern shown in fig. 1 can also explain the rather puzzling insensitivity of the fractionation factor for the proton in flight to the symmetry of the transition state found for instance in the diazo system. Some results for catalysis by L30+ are collected in the table. An advantage of using L30+ is that in all the cases except one one can also measure the degree of proton transfer (aL30+)from the secondary solvent isotope effect ; for diphenyl-diazo-methane we can see from fig.1 of Kreevoy's paper that a must be small. The results in table 4 of Kreevoy's paper show that for other acids than L30+ again -0.25 despite the large difference in reactivity TABLE O 1.-VALUES OF ~ L + ~AND dl FOR R1R2CN2 R1 RZ log(k/M-1 s-1) a aL30+ dl b IHC coo- 4.8 0.30 0.24 Me COOEt 1.3 0.29 0.22 CaSe C6H5 -0.1 t0.2 0.20 Me COMe -0.1 0.27 0.22 COOEtf COO- -1.7 0.36 0.27 a Rate constant for H+ catalysed reaction ; b Fractionation factor for proton in flight ; C M. M. Kreevoy and D. E. Konasewich J. Phys. Chem. 1970,74,4464 ; d W. J. Albery and A. N. Campbell-Crawford J.C.S. Perkiti ZI 1972 2190; e Ref.(1); f W. J. Albery C. W. Conway and J. A. Hail J.C.S. Perkin II 1976 473. and transition state symmetry for the 4 different acids. The same pattern and values of 41are found for other aliphatic diazo compounds.2 As discussed above this system has very large W'terms and quite small AGf terms. It corresponds therefore more closely to a very flat cirve B in fig. 1. This therefore means that there will be very little contribution from tunnelling to a maximum in the primary isotope effect (minimum in 41>.Furthermore the Westheimer contribution to the maximum will be small because first the fractionation factor for the proton after the W' process may be reduced from that in the reactant and second systems with small values of A require less unequal force constants to shift the Bronsted coefficients.Thus we may expect systems with large values of W' and small values of 1 not to show a pronounced maximum in the primary kinetic isotope effect as the symmetry of the transition state is changed. Prof. E. F. Caldin (University of Kent) said It is now recognised that the transfer of a proton and the resulting solvent reorganisation are not necessarily synchronous (cf. the papers by Bordwell Marcus Kreevoy et al. and Caldin and Wilson and ref. (23) of the last-named paper). For reactions where they are markedly out of step the general question arises which comes first ? Does the proton-transfer occur first and the resulting charge-separation force the polar solvent molecules to rotate ? Or does the proton-transfer occur only when the random motions of solvent molecules A.I. Hassid M. M. Kreevoy and T. Laing paper at this Symposium. W. J. Albery A. N. Campbell-Crawford and R. W. Stevenson J.C.S. Perkin II 1972 2198. F. €3. Westheimer Chem. Rev. 1961 61 265. GENERAL DISCUSSION lead to a configuration suitable to the transition state? Is there any general method of predicting the result or of determining it experimentally? This question is distinct from the question of coupling in which one process gives rise to a force that brings about the other. Rotation of solvent molecules takes a much longer time than the passage of a proton across an energy-barrier but could be coupled with it in the sense that the charge-separation associated with the proton- transfer could lead to a torque on polar solvent molecules which would produce rotation.It is also necessary to distinguish on the molecular level between processes that are “fast ” in the sense that the process requires a relatively short time (e.g. a vibration compared to a rotation) and those that are “fast ” in the sense that the process occurs frequently so that the associated rate is high. The rotation of solvent molecules though “slower ” than proton-transfer in the first sense may (or may not) be “faster ” in the second. Prof. R. A. Marcus (Uniuersity of Illinois) said In answer to Caldin’s interesting query regarding slow and fast coordinates I would like to present a picture or two which also serves to supplement Albery’s remarks.We consider first the nature of the potential energy profile. The profile depends of course on the coordinate being used as abscissa. For example if in the skewed-axis figure of my paper (fig. 1 there) the potential energy were plotted along the valley of the reactants over the saddle- point and along the valley of the products the barrier would be an Eckart-like one i.e. a roughly bell-shaped curve. If instead the potential energy were plotted against some collective solvent orientational coordinate allowing the proton to vibrate but not to transfer the profile would be as in fig. l(a)of this comment labelled reactants while if the protonic binding were that of the products the profile would instead be the curve labelled products in fig. l(a).The effect of a proton transfer on these profiles is indicated by the dotted-line profiles in fig. l(a) which allow for a protonic motion which always adjusts itself (adiabatically) to changes in the solvent coordinates rather than not being allowed to transfer. The splitting A&in fig. l(a) is related to the frequency of proton jumping v by the relation A&-hv,. When the proton jumps very readily (e.g. v -1014 s-l) A&is large perhaps roughly the size of a vibrational quantum for a hydrogenic motion. Fig. l(a) applies to a thermoneutral system. For a fairly exothermic reaction these profiles would instead be as in fig. l(b) while for a fairly endothermic one they would be as in fig. l(c). Thus turning to Caldin’s question in fig. l(b) where the U b C I FIG.1.-Profiles of potential energy for proton transfers versus some generalized solvation coordinate for the cases of (a) thermoneutral (b) exothermic and (c) endothermic reaction.The splitting AE reflects the proton jumping frequency -Ac/h at the cited value of this solvation coordinate. The proton jump occurs mid-way in largely prior to and largely after the solvation reorganization as indicated in cases (a),(b)and (c) respectively. S 10-6 142 GENERAL DISCUSSION “ intersection ” of the parabolas occurs early the protonic jump precedes the slow solvent molecules’ reorientations while in fig. l(c) the reverse is true. In fig. I(a) the slow reorientation partly precedes and partly follows the proton jump. Similar remarks apply when profiles are plotted against other slow coordinates.Incidentally in fig. 1 of this comment the protonic tunnelling is not a tunnelling through the barrier there. Rather it is reflected in the magnitude of v, i.e. in A&. Dr. W. J. Albery (Oxford University) said In considering the role of the solvent and the timing of its motion with respect to atom transfer I believe it is helpful to consider fig. 1 which shows a reaction involving proton transfer and a change in the solvation with a transition state that is intermediate on both co-ordinates (cf. fig. 1 p. 141 above).’ The effect of the torque exerted by the charge separation on the polar solvent molecules will be included in the calculation of the free energy surface and will be responsible in part for the shape of the profile along the proton transfer co-ordinate.The relaxation of the solvent after the proton transfer can be seen at the back of the diagram. Thus this type of “ coupling ” is described by the shape of the surface. As discussed earlier the most favourable path for the proton transfer will be when ApK for the actual proton transfer is not too far from zero. Movement along the solvation co-ordinate achieves this condition. However movement along this co-ordinate is slower than movement along the proton transfer co-ordinate. The translational and rotational diffusion of the solvent will have a lower velocity than the actual proton transfer. Note that one is concerned with the velocity of motion (cm s-l) rather than the velocity of reaction (M s-’).The latter is a function as well of the population of the different species. One could describe the system in terms of rate constants for motion along the different co-ordinates. However a more plausible model is to treat motion along the solvent co-ordinate as a diffusion process. Neutron diffraction data show that diffusion coefficients measured under macroscopic conditions continue to be a good description of motion down to times as short as 10-l2 s. We can write for the steady state where along the solvation co-ordinate AG --glx. RT In this equation k(x) is a function of the displacement along the solvation co- ordinate. If k(x) = 0 then there will be zero flux and one obtains the normal thermodynamic distribution for c c/c = exp( -g’x) = exp( -AGIRT).Analytical solutions for the flux can be obtained if one assumes that k(x) is zero for x < x1and then has a finite value. The form of the solutions their dependence on D or k,3 and whether the diffusion or the kinetic process is rate determining have been di~cussed.~ If the kinetic process is rate determining then the transition state is at + and the solvent is pre-equilibrated to the value at x,. If k(x) is large the diffusion process along x that is the solvent reorganization can be rate determining. W. J. Albery Reaction Transition States (ed. J. Dubois) (Gordon and Breach London 1972) p. 224. W. J. Albery Proton-Transfer Reactions (Chapman and Hall London 1975) p. 307. J. W. White Ber. Bunsenges. Phys.Chem. 1971 75 379. W. J. Albery Proton-Transfer Reactions (Chapman and Hall London 1975) p. 303 et seq. GENERAL DISCUSSION Solvation -FIG.1.-Diffusion/kinetic model of a proton transfer where the alteration of the solvent is shown by the square circle and diamond. The ApK of the AH S system varies as shown and proton transfer is most likely to occur when ApK = 0. The route through X where the solvent "lags " the proton transfer is less favourable than the route through f. Motion along the solution co-ordinate is assumed to obey a modified diffusion equation. Free energy differences corresponding to the Marcus parameters are shown except that the contribution to Wr from the formation of the encounter complex cannot be shown on the diagram.Solvation FIG.2.-Diffusion kinetic model for an asymmetric proton transfer where AG for the reaction is positive going from left to right and negative going from right to left. In the first case there is a large Wr term and solvent motion precedes the proton transfer. In the second case there is a small Wr term and solvent motion follows the Droton transfer. 144 GENERAL DISCUSSION Whatever the exact form of the solvation the flux can be described as the product of the concentration at x = 0 and a rate constant. Thus this analysis can be carried out for a succession of steps where the faster subsequent processes provide the rate constant "k " for the particular slower process being considered as a diffusion. Remember that by "faster " we do not niean the usual kineticist's distinction between fast and slow steps but rather we are concerned with the actual velocity of motion or the time taken for a single particle to complete that step of the reaction.In these terms the slower the process the earlier it has to take place. There will not be time for it to take place later as the particular molecule accelerates to its destruction. Thus for a typical reaction in solution we have the sequence Encounter + Ionic Atmosphere * Solvation + Atom transfer. Debye-Hiickel theory works for salt effects. Similarly there is no solvent lag because those reactants that are peculiarly solvated for the transition state are the ones that react. Reactants with normal reactant solvation do not react because for them the barrier through X (in fig.1) is larger than that through the transition state. The surface in fig. 1 is drawn for a reaction where AG = 0 and shows therefore a symmetrical transition state. For "uphill " transfers we obtain the surface shown in fig. 2. Here there has to be more solvent re-organization before ApK for the actual proton transfer becomes close to zero. Hence for the uphill case solvent re-organization largely precedes the proton transfer and W' is large. The "down-hill "case can be considered as the reverse of fig. 2. For this case there is little solvent re-organization before the proton transfers.

ISSN:0301-5696    

DOI:10.1039/FS9751000132

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Proton Bridges in Enzyme Catalysis BY J. P. ELROD,R. D. GANDOUR, J. L. HOGG,M. KISE G. M. MAGGIORA AND K. s. VENKATASUBBAN R.L. SCHOWEN* Departments of Chemistry and Biochemistry University of Kansas Lawrence Kansas 66045 U.S.A. Received 28th April 1975 Chains of two hydrogen bonds (charge-relay systems) exist in the active sites of serine proteases and may be catalytic entities. It is shown theoretically that the coupling of the motions of the two protons in such arrays (and thus the efficiency with which they can relay charge) is a critical function of the distances across the hydrogen bonds ; long distances favour uncoupled motion short distances coupled motion. Rate measurements in mixtures of light and heavy water show that the serine proteases chymotrypsin trypsin and elastase function as one-proton catalysts one proton of the chain presumably bridging as in ordinary general catalysis.On the other hand three enzymes of the amidohydrolase class a glutaminase and two asparaginases show two-proton catalysis. This may arise from a charge-relay chain although evidence for such a structure has not yet been advanced for these enzymes or from some other catalytic entity involving two coupled proton bridges. Thirty-four years ago R. P. Bell’s Acid-Base Catalysis laid out the foundations on which he and several generations of scientists have build today’s edifice of under- standing in this field. While the importance of acid-base interactions for the catalytic power of enzymes has been clear for a long time,’ the recent growth of information on the three-dimensional structure of enzymes has led to the discovery of active-site entities of particular catalytic potential.Among these is the charge-relay chain (eqn (I)) found in the active sites of such serine proteases as chymotrypsin trypsin and elastase and recognized by Blow Birktoft and Hartley as a possible source of catalysis. As negative charge is relayed to the hydroxyl oxygen as in eqn (l) its (asp) (his) (ser) (asp) (his) (ser) nucleophilicity is increased and increased to an unusual extent because the protonic positive charge is transferred out to the carboxyl group rather than merely to the neighbouring imidazole. Conversely relay of the proton (reverse of eqn (1)) to the hydroxyl oxygen (or to another group occupying its position) can greatly increase the leaving-group reactivity of this centre because the negative charge is removed to a correspondingly large distance.Charge-relay chains should be encountered only rarely in non-enzymic systems (excepting the special case of solvent bridging) because of the entropic difficulties of assembling and orienting the composite group^.^ To function most effectively in the manner just described the charge-relay chain must require some degree of coupZing of the motions of the two protons in its com- ponent hydrogen bonds. In this way a smooth relay of charge is accomplished while only an ordinary catalytic advantage would be derived from independent shifts of the protons. In this paper we describe a theoretical demonstration of the impor- 145 PROTON BRIDGES tance of geometry for the coupling of proton motions in hydrogen-bond chains and experimental tests of the degree of coupling in some enzymic activated complexes which potentially involve multiple-proton bridges.COUPLING OF PROTON MOTIONS IN HYDROGEN-BOND CHAINS In an array such as the charge-relay chain the potential energy E may be described as a function of two variables Arl and Ar2 representing the displacements of the two protons from their equilibrium positions with all other coordinates of the system being either relaxed to their minimum-energy values or constrained to some desired values (e.g. to the known positions of groups in an enzyme active site). This permits construction of a three-dimensional potential-energy surface E(Arl Arz) and its representation by a contour map of the familiar sort.In discussing coupling of the two motions it is convenient to plot energy as a function of normalized coordinates p1 = (Arl/Art) and p2 = (Ar2/Ari),where Arf represents the total displacement of the proton which occurs in the overall reaction. The coordinates p1 and p2 are dimensionless quantities which vary from zero to one as the reaction proceeds from reactants to products. Fig. 1 illustrates a schematic contour map E(pl,p2) and shows I I / c / I I / / I / 0 P2 FIG.1 .-Schematic potential-energy surface E(pl pz) in which energy contours would be plotted as functions of the normalized coordinates pl = Ar,/Ari and pz = Two possible pathways representing the longest (L = 2) and shortest (L = 2+),are shown.two possible minimum-energy reaction paths from reactants (0 0) to products (1 1). The diagonal pathway represents the most direct way of accomplishing the reaction with both proton motions proceeding in perfect synchrony (thus exactly coupled). The dimensionless length of this pathway (projected onto the plane E = 0) is 2*. The other pathway represents the most indirect reaction scheme-a perfectly stepwise or uncoupled path. Its dimensionless length (projected on E = 0) is 2. Thus the length L of the projected reaction path across E (pl p2) is a quantitative measure of the overall (dynamical coupling during reaction. In fact we can define a degree of coupling oby eqn (2) so that ovaries from zero for perfectly uncoupfed processes to unity for perfectly coupled processes.(The extension to other processes and to spaces of higher dimensionality is straightforward with the limits of L being N and N3 where N is the number of coordinates 5b) 0= (2-L)/(2-23). (2) ELROD GANDOUR HOGG KISE MAGGIORA SCHOWEN VENKATASUBBAN 147 DETERMINANTS OF COUPLING IN PROTON BRIDGES It seems initially reasonable that the chemical constitution of a hydrogen-bond chain ought to influence the degree of coupling exhibited by a hydrogen-bond chain; this has been confirmed the~retically.~~ The same study showed that a second factor of great and perhaps dominant importance was the distance across each hydrogen bond. Long distances favoured uncoupled reactions while shorter distances led to ever greater coupling.This finding may hold particular significance for the hydrogen- bond arrays of enzymes where the distances are effectively fixed by the three- dimensional structure (thus by molecular evolution) but may also be to some degree adjustable for example by the binding of substrate or through conformation changes of the enzyme. rFH FIG.2.-Potentialenergy surface for the conversion of +H3N-H.. .F-H.. .OHz (lower left corner) to H3N...H-F.. .H-OH; (upper right corner). Energy contours are labelled in kcal/mol. The N-F and F-0 distances are constrained to 3.00 A. The ordinate is ~NH and abscissa is ~FH. Since A&H = 0.90 8,and A& = 0.96 A a rough transformation to normalized coordinates was performed by setting ArhH-ArkHd.93 A.Then L = 1.74 (w = 0.44) for the “ northwest ” reaction path and L = 1.77 (w = 0.39) for the “southeast ” reaction path. Calculations to illustrate the effect of hydrogen-bond length on coupling are shown in fig. 2 and 3. Both calculations are for the reaction of eqn (3) with the distance R held constant at 3.00A in fig. 2 and at 2.75A in fig. 3. A discussion of the computational method and the basis for choice of the chemical constituents in eqn (3) can be found in an earlier publi~ation.~~ As is clearly seen in fig. 2 and 3 the surfaces change character completely when R is altered by only 0.25 A. At the longer distance seen in fig. 2 there are two minimum energy paths (“ northwest” and “ southeast ”) with low degrees of coupling.For the northwest route o = 0.44 and for the southeast o = 0.39. On the other hand when the system is made shorter as in fig. 3 there is a single highly coupled path (o= 0.94). Thus compres- sion of the chain converts a very poorly coupled system into one which is nearly perfectly coupled. PROTON BRIDGES tR+cR+ tR+tR+ +HSN-H.. .F-H.. .OH,+ H3N...H-F...H-OHl (3) This behaviour is entirely consistent with expectations from the component hydrogen-bond potential functions. A long hydrogen bond should have a high barrier between its two potential-energy minima because a centrally-located proton will be weakly bound by both bases. In the shorter system acentral proton can occupy the overlapping region of the potential fields of the two bases decreasing the barrier.Coupled motion of the two protons in a chain leads these effects for the individual hydrogen bond to reinforce producing a high central region for the surface of a long system and a lower central region for the shorter system. rFH FIG.3.-PotentiaLenergy surface for the conversion of +H3N-H.. .F-H.. .OHz to H3N...H-F.. . H-OH; with N-F and N-0 distances constrained to 2.75A. Energy contours are labelled in kcal/mol. The ordinate is ~NHand the abscissa ~FH. Since A~&H= 0.56A,arb^ = 0.61 A a rough transformation to normalized coordinates was made with Arh~-Arb~-0.58A. This yields L = 1.45 for the reaction path and w = 0.94. PROTON INVENTORIES OF ENZYMIC TRANSITION STATES It is also desirable to investigate experimentally the degree to which coupled motion of the protons in enzymic charge-relay chains contributes to catalysis.Since each proton should produce a kinetic isotope effect if its binding state is altered on formation of the catalytic transition state the information wanted is a list of" active protons" (those generating isotope effects) and the isotope effect for each a "proton inventory ". Such an inventory is in principle accessible by studies in mixtures of protium and deuterium oxides?. According to the well-established relationship of eqn (4) k, the rate constant in a solvent mixture with atom fraction n of deuterium is related to ko (for pure protium oxide) through the reactant-state isotopic fractionation factors 47 and the transition-state isotopic fractionation factors 4:.If as eqn (4b) ELROD GANDOUR HOGG KISE MAGGIORA SCHOWEN,VENKATASUBBAN 149 emphasizes the entire reactant-state contribution RSC(n) is known the relationship becomes a polynomial in n the order of which (v) specifies k = kon(1-n+n4T) 1 V k,RSC(n) = ko n(1-n +n#T) (4b) i the number of " active " protons and the coefficients of which allow the calculation of the +T,i.e. the isotope effect for each proton. Thus k,(n) along with some measure or hypothesis concerning RSC (n),can yield the proton inventory. Since only those protons which change state on activation will contribute to RSC(n) and since exchangeable protons of proteins (except those of sulphydryl groups) are expected to have 4 -1 we set RSC(n) equal to unity for the cases to be considered here.There can be some dangers in this procedure. Kresge has shown that small reactant contributions may conspire under cover of strong coincidence and experimental error to conceal a transition-state contribution to k,(n). Although it is desirable to be alert to this possibility it is unlikely that Nature will haunt a system from case to case producing exactly the appropriate cancellation conditions in each event. Thus a similar result with different substrates or under different conditions ought generally to lay this ghost. For the particular case of double-hydrogen-bond chains such as the charge-relay chain one can distinguish two extreme situations and a continuum of intermediate cases.If the system is totally uncoupled but with one of its components still acting as a one-proton catalytic bridge (v = 1 in eqn (4b),analogously to common general catalysis lo) then k,,(n)should be linear in tz. If the system is perfectly coupled with both hydrogens generating equal kinetic isotope effects (v = 2 4; = $;) then k,(n) should be quadratic and [k,(n)]%should be linear in n. Between these extremes there should be a continuum of imperfectly coupled cases in which k,(n) is quadratic but 4; # 6;. These and the perfectly coupled case correspond to the two-proton catalytic bridge or true charge-relay. ONE-PROTON BRIDGES IN SERINE PROTEASES Proton inventories for two reactions of a-chymotrypsin are shown in fig. 4. The filled circles in fig.4 are data for removal of an acetyl group attached to the serine hydroxyl of the charge-relay chain. The measurements are those of Pollock Hogg and Schowen," enriched by some later points obtained by Elrod. The dependence of k on n is clearly linear corresponding to a single-proton bridge with isotope effect kH/kD = 2.4. Two objections may be raised to this experiment. First coincidental cancellation of factors from RSC(n) with the contribution of a second trmsition- state proton can conceal its effe~t.~ Second the deacetylation of the acetyl-enzyme is rather remote from the physiological process for which evolution designed chymo- trypsin and therefore even if the charge-relay chain is uncoupled in this reaction it is possible that the full structure of a physiological substrate would call the coupled chain into action.Both objections are met by the open circles of fig. 4 for the acylation of the enzyme (nucleophilic attack by the serine hydroxyl) by a close analog of the natural peptide substrates N-acetyl-L-tryptophanamide (ATA). Since this quite different process again produces a hear k,(n) with kH/kDnow 1.9 it is most unlikely that fortuitous cancellations are at work. Further this substrate fills all the binding PROTON BRIDGES n FIG.4.-Ratio of maximum velocity Vn in mixed isotopic solvent to Vl for deuterium oxide as a function of n for deacetylation of acetyl-a-chymotrypsin (25.00k 0.05" pH 7.5 and equivalent filled circles) and for acylation of a-chymotrypsin by N-acetyl-L tryptophanamide (ATA 25.00 & 0.5" ; pH 8.10 and equivalent open circles).positions of the active site and ought to simulate well the physiological situation. Nevertheless the charge-relay chain appears uncoupled and the enzyme is employing a one-proton catalytic bridge. Studies concerned with two other members of the serine-protease family appear in fig. 5. The open circles of fig. 5 are for the deacetylation of acetyltrypsin (quite analogous to the chymotrypsin reaction of fig. 4) which exhibits a linear k,(n) with n FIG.5.-Ratio of maximum velocity V in niixed isotopic solvent to Vl for deuterium oxide as a function of n for deacetylation of acetyltrypsin (25.00+0.10" pH 7.54 and equivalent open circles) and of acetylelastase (25.Ook 0.10"; pH 7.54 and equivalent filled circles).ELROD GANDOUR HOGG KISE MAGGIORA SCHOWEN VENKATASUBBAN 151 kH/kD= 1.4. Again one would prefer to add results for a physiological substrate but our studies in this regard are incomplete. However although details have not been published Mason and Ghiron l2 have reported the rates of deacylation of N-benzoyl-L-arginyltrypsinin mixed isotopic solvents. Their values for k,(n) generate a linear plot (not shown) with kH/k,= 2.6. Therefore it seems clear that trypsin also employs a one-proton catalytic bridge. Finally the results for deacetyl- ation of acetylelastase a third member of the family seen as the filled circles of fig. 5 are also indicative of the one-proton catalytic bridge. The isotope effect is kH/kD= 2.2.Results for a physiological substrate of elastase are not yet in hand. The serine proteases seem not to employ charge-relay as a catalytic factor but rather a one-proton bridge of the type familiar in general catalysis. The small magnitudes of the isotope effects (k,/k = 1.4-2.6) suggest to us that the proton in these bridges is not "in flight " but is a non-reaction-coordinate proton engaged in "solvation catalysis ''.lo? l3 TWO-PROTON CATALYSIS BY AMIDOHYDROLASES Other enzymes in addition to the serine proteases have as their physiological task the catalytic hydrolysis of the amide linkage. Among these are the aspara- ginases and glutamina~es,~ which catalyze the hydrolysis of asparagine (eqn (5) x = 1) and glutamine (eqn (9,x = 2) respectively.These enzymes are larger and more complex than the serine proteases detailed X-ray crystallographic structures are not available and mechanistic investigations are at a more primitive stage. Thus we have no indication as yet whether a charge-relay chain exists at the active sites of l-i'l \ GLUTAMINASE 1.01 L \x ASPARAGINASE, 1 1 ERWIN;\ 0 0.5 1 n FIG.6.-Square root of the ratio of maximum velocity V, in mixed isotopic solvent to VI for deuterium oxide as a function of n for amidohydrolase reactions. The upper plot is for glutamine hydrolysis catalyzed by glutaminase of E. coli at 37.0*0.2" pH 5.50 and equivalent. The lower plot is for hydrolysis of asparagine by asparaginase of E. coli (37.00+0.2" pH 7.12 and equivalent open circles) and by asparaginase of Erwinia carotoooru (37.00i-0.02" pH 7.12 and equivalent filled circles).PROTON BRIDGES these enzymes. Nevertheless it seems that like the serine proteases they form an intermediate acyl-enzyme and it is of interest to inquire to what extent they share the same proton-bridging properties in their catalytic transition states. -0,c 0 -0,c \ // \ -b CH(CH2)XC CH(CH,),CO +NH (5) / \ / HiN NH H;N Fig. 6 displays results from three enzymes a glutaminase of Escherichia coli an asparaginase of Escherichia coli and an asparaginase of Erwinia carotovora. All are catalyzing the hydrolysis of their natural substrates glutamine or asparagine at 37”. In this figure the square-root [kn(n)]*is plotted and is a linear function in all three cases.Thus all three enzymes are employing two-proton catalytic bridges with k,/k = 1.33 (glutaminase) 1.71 (asparaginase E. coli) and 1.62 (asparaginase Erwinia carotouora) for each proton and thus overall solvent isotope effects of 1.80 2.93 and 2.62 respectively. It then appears that in the amidohydrolases two protons are coupled in the catalytic transition state. The structural nature of the catalytic bridge(s) remains unknown a charge-relay chain a bifunctional catalytic entity or some other apparatus may be at work. It is tempting to speculate that a charge-relay chain exists here and that it has in these enzymes been brought to such an overall length through structural factors that its proton motions are now closely coupled.APPENDIX EXPERIMENTAL DETAILS The INDO potential surfaces for the asymmetric hydrogen-bond chains +H3N-H.. . F-H.. .OH2,were calculated in the manner described for symmetrical systems by Gandour Maggiora and Sch~wen.~~ Rate measurements were generally made spectrophotometrically employing a Cary 16 spectrophotometer interfaced to a Hewlett-Packard 2100A computer. The output of the photomultiplier consisting of a 60 Hz train of alternating sample and reference pulses was conveyed to an analogue-to-digital converter through a synchronizer which identified the pulses. Fifteen measurements of the height of both sample and reference pulses were averaged across each cycle and a value of the absorbance calculated from the logarithm of the ratio of these averages.Reaction times were divided into 1000 segments and absorbances were then time-averaged across each segment. Zero-order rate constants were calculated from a linear least-squares fit of absorbance to time. For some glutaminase runs ammonium-ion concentrations were determined by a Beckman cation-sensitive electrode (39137) interfaced to the same computer. Enzymes with one exception were commercial materials obtained from Sigma Chemical Company [chymotrypsin trypsin elastase glutaminase and asparaginase (E.coli)] or Worthington Biochemicals (chymotrypsin). The asparaginase of Erwinia carotovora was a generous gift of Dr. H. E. Wade of the Microbiological Research Establishment Porton Down Salisbury U.K. Deuterium oxide was obtained from various sources (Diaprep Biorad and Stohler Isotope Companies) and was either distilled or used as supplied; in all cases rates were shown to be independent of the method of purification.Tris buffer components were also obtained from Sigma and were used for pH control near pH 8. Acetate buffers were used near pH 5. Constant buffer ratios in mixed isotopic solvents were maintained in order to hold the pL at a corresponding point on the pL/rate profiles (an “ equivalent pH ”). In a typical experiment 3 ml of buffer solution was placed in a cuvet in the thermostatted cell-holder of the spectrophotometer and allowed to attain thermal equilibrium with the ELROD GANDOUR HOGG KISE MAGGIORA SCHOWEN VENKATASUBBAN 153 bath as indicated by a glass-covered thermistor probe.Then 0.1 ml of the enzyme stock solution was introduced by micropipet. After five minutes 0.1 ml of substrate stock solution was injected the solution was shaken and the absorbance monitored at 215 nm (amide substrates) or 400 nm (p-nitrophenyl acetate). It is a pleasure to acknowledge the support of this work by the National Institutes of Health and National Science Foundation (U.S.A.) the kind gift of Erwinia asparaginase from Dr. H. E. Wade of the Microbiological Research Establishment Porton Down Salisbury (U.K.) and a grant of computer time from the University of Kansas. W. P. Jencks Catalysis in Chemistry and Enzymology (McGraw-Hill Book Co. New York 1969) ; M. L. Bender Mechanisms of Homogeneous Catalysis from Protons to Proteins (Wiley-Interscience New York 1971).D. M. Blow The Enzymes ed. P. D. Boyer (Academic Press New York 3rd edn. 1971). vol. 3 chap. 6; G. P. Hess chap. 7; B. Keil chap. 8; B. S. Hartley and D. M. Shotton chap. 10. D. M. Blow J. J. Birktoft and B. S. Hartley Nature 1969 221 337. 6.A. Rogers and T. C. Bruice J. Amer. Chem. Soc. 1974 96; 2473 ; M. Choi and E. R. Thornton J. Amer. Chem. SOC.,1974 96 1428. (a)R. D. Gandour G. M. Maggiora and R. L. Schowen J. Amer. Chem. SOC. 1974,96,6967. (b) P. Hogan R. D. Gandour G. M. Maggiora and R. L. Schowen to be published. A. J. Kresge Pure Appl. Chem. 1964 8 243. V. Gold Adv. Phys. Org. Chem. 1969 7 259. R. L. Schowen Progr. Phys. Org. Chem. 1972 9 275. A. J. Kresge J. Amer. Chem. SOC. 1973 95 3065.lo S. S. Minor and R. L. Schowen J. Amer. Chem. SOC.,1973,95,2279. E. Pollock J. L. Hogg and R. L. Schowen J. Amer. Chem. Soc. 1973 95 968. '* R. Mason and C. A. Ghiron Biochim. Biophys. Acta 1961 51 377. l3 C. G. Swain D. A. Kuhn and R. L. Schowen J. Amer. Chem. Soc. 1965 87 1553. Is J. C. Wriston Jr. The Enzymes ed. P. D. Boyer (Academic Press New York 3rd edn. 1971) vol. 4 chap. 5 ; J. C. Wriston Jr. Adv. Enzymol. 1973 39 185. l5 S. C. Hartman The Enzymes ed. P. D. Boyer (Academic Press New York 3rd edn. 1971) vol. 4 chap. 4.

ISSN:0301-5696    

DOI:10.1039/FS9751000145

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

The Nature of the Proton Transfer from an Acid Group at the Active Site of an Enzyme to Solvent Water The Extent of 2Hand 3HTransfer in the Reaction catalysed by Triose Phosphate Isomerase BY L. MARKFISHER,* w. JOHNALBERY? AND JEREMYR.KNOWLES" Department of Chemistry Harvard University Cambridge Mass.( *) and Physical Chemistry Laboratory University of Oxford(?) Received 2nd May 1975 The extent of transfer of a 2H and of a 3H label in 1-R-[*H or 3H]-dihydroxyacetone phosphate to the 2-position of D-glyceraldehyde 3-phosphate in the reaction catalysed by triose phosphate isomerase has been determined. Since the enzymic base that abstracts the substrate's isotopic label is a carboxylate group this enzyme-substrate system effectively provides -COOZH and -CO03H in a solution of 'H2O,and allows an investigation of the transfer of ZH and 3H from the carboxyl group to unlabelled water.From the value of the fractionation factor for this proton transfer it is evident that the mechanism of exchange of isotope on -COOL with the protons of the solvent does not involve a transition state in which L is in flight. The glycolytic enzyme triose phosphate isomerase catalyses the interconversion of dihydroxyacetone phosphate (I) and D-glyceraldehyde 3-phosphate (IIj through abstraction of the 1 pro-R hydrogen of the dihydroxyacetone phosphate or of the hydrogen at C-2 of the D-glyceraldehyde 3-phosphate by a catalytic base at the active site of the enzyme.l Experiments with specific active-site-directed inhibitors and recent crystallo- graphic work on the enzyme strongly suggest that this base is the y-carboxyl group of a glutamate residue.This assignment is consistent with the pH-dependence of the enzyme activity and attractively accommodates a number of other features of the catalysed reaction. When 1-R-[3H]-dihydroxyacetonephosphate is incubated with isomerase under conditions where the initial product D-glyceraldehyde 3-1 54 L. M. FISHER W.J. ALBERY AND J. R. KNOWLES phosphate is oxidised to 3-phosphoglycerate as fast as it is formed only some 6% of the tritium label originally at C-1 is transferred (after complete reaction) to C-2 of the final product 3-phosphoglycerate. This demonstrates that the catalytic base having abstracted the carbon-bound proton can exchange this proton with the solvent water during the course of the enzyme-catalysed reaction.A plausible scheme for catalysis by the isomerase involving an enediol intermediate has been proposed (see fig. I) and a generalised kinetic treatment has been derived that allows the rate constants for the individual steps of the reaction to be evaluated from the results of a number of isotopic experimenk6 This paper is concerned with the isotope exchange reaction between the conjugate acid of the catalytic base B-L (fig. l) and solvent FIG.1.-Transfer of hydrogen isotope in the isomerisation of dihydroxyacetone phosphate by triose phosphate isomerase. B-is the active-site carboxyl group L is 2Hor 3H,@ represents a phosphoryl group and the heavy arrows indicate the fate of the isotopic label.All the species represented are enzyme-bound. water. The extent of deuterium transfer from 1-R-[2H]-dihydroxyacetonephosphate to the final product 3-phosphoglycerate has been measured under conditions essentially identical to those of the tritium transfer experiment mentioned above.5 The results from these two experiments are used to derive the isotope effect for the reaction involving proton exchange between solvent water and the carboxyl group at the active site of the isomerase. The enzyme-substrate system effectively provides -C002H and -CO03H in a solution of 'H20 and allows an investigation of the transfer of 2Hand of 3H from the carboxyl group to unlabelled water. From the full kinetic treatment of the reaction scheme laid out in fig.1 the fraction of isotope transferred (p") from 2H- or 3H-labelled I -R-dihydroxyacetone phosphate to the 2-position of the product 3-phosphoglycerate under conditions where the isomerase reaction is rate-limiting approximates to where k refers to the rate of the irreversible isotope exchange reaction with solvent ks B-L + H2O + B-H + HOL and kt is the composite rate constant for the collapse of the enediol to D-glyceraldehyde phosphate and its departure from the enzyme's active site. From the knowledge of the partitioning ratio of the labelled enzyme-bound enediol intermediate (i.e. the extent of transfer of 2H and of 3H) the isotope effect for the k exchange reaction of eqn (2) is obtained from the Swain-Schaad equation.' PROTON TRANSFER TO WATER FROM AN ENZYME ACID EXPERIMENTAL Dihydroxyacetone phosphate stereospecifically labelled with deuterium in the 1-R position was prepared by equilibration of dihydroxyacetone phosphate in 2Hz0 (99.8 %) using chicken muscle triose phosphate isomerase.8 Ion exchange chromatography yielded l-R-[2H]-dihydroxyacetonephosphate essentially free of ~-2-[~H]-glyceraldehyde 3-phosphate (at equilibrium the mixture contains 96% of dihydroxyacetone phosphate and 4% of D-glyceraldehyde phosphate) and the extent of deuterium incorporation was determined by mass spectrometric analysis of the volatile derivative 1-[2H]-tetrakis(trimethyIsilyl)-a-glycerolphosphate using an AEI MS9 instrument.For the transfer experiments 1-R- [2H]-dihydroxyacetone phosphate NAD+ EDTA and sodium arsenate were dissolved in 100 mM triethanolamine-HC1 buffer pH 7.6 and equilibrated in an optical cuvette at 30°C.The amount of any contaminating ~-2-[~H]-glyceraldehyde 3-phosphate was determined by enzymic assay using isomerase-free glyceraldehyde 3-phosphate dehydrogenase. The isomerase reaction was then initiated in the same cuvette by the addition of a small (rate- limiting) quantity of triose phosphate isomerase. The progress of the reaction was foilowed to completion by monitoring the increase in absorbance at 366 nm due to the formation of NADH. The product of the reaction 3-phosphoglycerate was isolated and purified by ion exchange chromatography. This material was methylated using diazomethane.Mass spectra of the methylated product were obtained at a probe temperature of 45" and 70 eV by direct insertion using an AEI MS9 instrument. The peaks at mle 169 and mle 170 were scanned slowly and repeatedly and the extent of deuterium transfer in the isomerase- catalysed reaction was determined from the averaged mle 170:169 intensity ratios of the labelled methylated 3-phosphoglycerate and of the unZabeZZed material obtained similarly. As a check of both the extent of labelling of the original dihydroxyacetone phosphate and of the position of the isotopic label in the product 3-phosphoglycerate a very little tritiated water was added to the 2H20 used for the labelling of the dihydroxyacetone phosphate. This radioactive tracer was also used (see below) as a check on the fate of the deuterium label.RESULTS AND DISCUSSION The intense (M+-59) peak at m/e 169 in the mass spectrum of methylated 3- phosphoglycerate is formed from (111) by facile loss of -COOCH,. Incorporation of deuterium at the C-2 position of 3-phosphoglycerate increased the intensity at m/e 170 relative to that at m/e 169. Data from the mass spectra of four samples of L \ OCH 3-phosphoglycerate obtained from separate deuterium-transfer experiments gave the following per cent deuterium contents 6.5 6.3 5.0 and 5.9 %. Each of these values was obtained from between 5 and 14 scans of the mass spectrum. The starting 1-R-[2Hj-dihydroxyacetone phosphate was shown by mass spectrometry to be essentially completely deuterated at C-1 and the percentage deuterium transfer was L.M. FISHER W. J. ALBERY AND J. R. KNOWLES corrected in each case for any small contribution from contaminating D-2-L2 HI-glyceraldehyde 3-phosphate. The extent of deuterium transfer from 1-R-[2H]-dihydroxyacetone phosphate to 3-phosphoglycerate catalysed by the isomerase is thus about 5.9%. To confirm the location of the isotopic label in the product 3-phosphoglycerate use was made of the tritium tracer initially incorporated into the starting 1-R-[2H]-dihydroxyacetone phosphate. A portion of the 3-phospho- glycerate product was treated with the enzymes phosphoglycerate mutase and enolase and the tritium specifically labilised from C-2 of the 3-phosphoglycerate (see fig. 2) was recovered as tritiated water by distillation.coo-COO-coo-\ L-C-OH L-C-O-@ \ I yLI I/ + CH2 CH2 CH LOH 0-@ OH FIG.2.-Check on the position of the isotopic label in 3-phosphoglycerate. a phosphoglycerate mutasei-2 3-diphosphoglycerate as cofactor ; b enolase. The isotopically-labelled conjugate acid of the base at the active site of the isomerase may suffer one of two fates (see fig. I). The label may be transferred to the enediol intermediate forming labelled D-glyceraldehyde 3-phosphate which is then lost from the enzyme and converted to 3-phosphoglycerate labelled at C-2 (the k route). Alternatively the label may be irreversibly exchanged for a proton from solvent water resulting ultimately in the formation of unlabelled 3-phosphoglycerate (the k route). Clearly the extent of transfer of label to C-2 of 3-phosphoglycerate is a measure of the relative fluxes of material along these two pathways.From the work of Herlihy et a/.,5we know that for tritium transfer pT = 0.058 f. 0.004. That is after complete conversion of 1-R-[ 3H]-dihydroxyacetone phosphate to product 5.8f0.4% of the tritium label is transferred intramolecularly. In the present work we have found that for deuterium transfer pD = 0.059+0.005. From eqn (1) we have where P is derived from the experimental values of pT and pD. From the Swain Schaad relationship for any reaction kD/kH= (kT/kD)2*3, so In terms of the deuterium fractionation factors & and 4t for the transition states for the two routes 4 = 4 P-2*3. (5) Now by performing the enzyme-catalysed reaction in tritiated water lo and comparing the tritium content of the product with that of the solvent we know the tritium fractionation factor on the k route Qt = 0.83f0.01 from which the Swain- Schaad relationship gives q5t = 0.88 f.O.01.This fractionation factor is slightly less than unity since the rate of transfer is controlled mainly by the loss of product from the enzyme (4 = 1.0) but partly by the proton transfer step as the bound enediol collapses to bound D-glyceraldehyde 3-phosphate (4-0.2 to 0.3). From the above PROTON TRANSFER TO WATER FROM AN ENZYME ACID equations we find the values of +s from our four experiments to be 1.1 1.1 0.61 and 0.92. +s is therefore 0.94f0.17. This value is close enough to unity for it to be clear that the mechanism of exchange of the isotope on -COOL with the protons of the solvent does not involve a transition state in which L is in flight.We may envisage three kinds of pathway for this isotopic exchange. First the pathway may involve ionisation solvent exchange and reprotonation (fig. 3). At pH 7 the solvent exchange step will be much faster than the diffusive encounter of H30+and B- and provided that the diffusive step is rate-limiting our observed value for the fractionation factor is consistent with this mechanism. diffusion -1.0 Jt solvent B-..*H-O + H,O+ reorgoniration far? B-... L -0 + H,O+ I 4 -1.0 H H FIG.3.-Stepwise proton exchange process between a labelled acid B-L and HzO. The exchange of a carboxyl proton -COOL however may occur in a cyclic manner within a hydrogen-bonded complex such as has been proposed to account for the H n.m.r.of acetic acid-acetate buffers (see fig. 4). In such a mechanism H /” L transfer H -0. c ‘H -R-C Lo-c ...0 ’ 4-0.2 ‘H H transfer L -4 -1.0 fast solvent I reorganisation FIG.4.-Cyclic proton exchange for a carboxylic acid the second proton transfer must be rate-determining since H20 is a weaker base than R-COO- and solvent reorganisation will be faster (at 108-109 s-l) than the cyclic proton transfer (106-107 s-l). The effect of L on the rate-determining transition state is now secondary and since the transition state is symmetrical the fractionation factor for this step will be ,/l = 40.69 = 0.83 (see ref.(12)) [where I is the fraction- ation factor for the process L2HO++3D20 +L2DO+++H20(ignoring statistical factors)]. This is also consistent with our experimental value. Since the cyclic mechanism is some 102-fold faster than the normal dissociation path we may expect L. M. FISHER W. J. ALBERY AND J. R. KNOWLES this route to be more probable provided that the steric constraints of the active site allow it. Thirdly it is conceivable that during the lifetime of the enediol intermediate (i.e. of B-L) there is complete and rapid exchange off with a limited pool of about 3 water molecules that are isolated from the bulk water during the reaction. In this case as soon as the substrate or product leaves the active site these water molecules would be free to exchange with bulk solvent and a fresh set of unlabelled water molecules would be trapped by the next substrate to bind.This situation could explain the absence of a measurable difference between the extent of transfer of 2H and 3H. We do not favour this interpretation since it does not agree with the characteristics of enzyme active sites in general nor with the active site of triose phosphate isomerase in partic~lar.~ From the crystal structure of the isomerase at high res~lution,~ it seems unlikely that the substrate could entrap 3 water molecules at the active site though this point will be completely clarified when the crystal structure of the isomerase-dihydroxyacetonephosphate complex has been solved at high resolution.In summary it is clear that a process not involving proton transfer from -COOL is the rate determining step for the overall stepwise proton exchange reaction. Analogous situations have been proposed for amines l3 and for a number of carbon acids,14 but this is the first analysis of the nature of the ionisation of a carboxylic acid in aqueous solution. S. V. Rieder and I. A. Rose J. Biol. Chem. 1959 234 1007. * S. de la Mare A. F. W. Coulson J. R. Knowles J. D. Priddle and R. E. Offord Biochem. J. 1972 129 321 ; S. G. Waley J. C. Miller I. A. Rose and E. L. O’Connell Nature 1970,227 181; F. C. Hartman Biochem. 1971 10 146. D. W. Banner A. C. Bloomer G. A. Petsko D. C. Phillips C. I. Pogson and I. A. Wilson Nature 1975 255 609. B. Plaut and J.R. Knowles Biochem. J. 1972 129 31 1. J. M. Herlihy W. J. Albery and J. R. Knowles to be published. W. J. Albery and J. R. Knowles to be published. For a preliminary description of some of the isotope experiments see J. R. Knowles P. F. Leadlay and S.G. Maister CoZd Spring Harbor Symposia on Quantitatice Biology 1971 36 157. ’C. G. Swain E. C. Strivers J. F. Reuwer Jr. and L. J. Schaad J. Amer. Chem. SOC.,1958 80 5885. * S. J. Putman A. F. W. Coulson I. R. T. Farley B. Riddleston and J. R. Knowles Biochem. J. 1972 129 301. T. Curstedt Eur. J. Biochem. 1974 49 355. lo S. G. Maister C. P. Pett W. J. Albery and J. R. Knowles to be published. Z. Luz and S. Meiboom J. Amer. Chem. SOC.,1963,85,3923. V. Gold Trans. Faraday Soc. 1968 64,2770; A. J. Kresge Pure Appl. Chem. 1964 8 243. l3 E. Grunwald and E. K. Ralph Acc. Chem. Res. 1971 4 107. l4 e.g. D. J. Cram D. A. Scott and W. D. Nielsen J. Amer. Chem. SOC.,1961,83 3696.

ISSN:0301-5696    

DOI:10.1039/FS9751000154

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

GENERAL DISCUSSION Dr. W. J. Albery (Oxford University) said In using the solvent isotope curve to draw conclusions about the transition state one has to be very careful. Fig. 1 shows a typical solvent isotope curve. Information is contained in the overall magnitude of the effect (the value at x = 1) and in the shape of the curve (the value at x = 3). Unfortunately all curves that pass through the same values at x = + and x = 1 are so close to each other at x = t and x = 3 that even rate ratios precise to several parts in a thousand cannot distinguish between the different curves corresponding to different models of the transition state. Hence it is better to concentrate experi- ments at x = 3 and x = 1. The extra information obtained for experiments at x = 3can be described by the curvature parameter which describes how far the point at x = 3 lies off the line joining the points at x = 0 and x = 1 as shown in fig.1 ; y is normalised so that y = 0 for the straight line A in fig. 1 y = 1 for the curve B which corresponds to a linear variation of k,/ko with x and y = + for curve C the " quadratic " case discussed by Schowen. When one has obtained a value of y instead of carrying out an " inventory of protons " one can carry out an inventory of the possible transition states which will fit the data and which have expressions of the form (1 -x+ (1 -x +&x)~. In this expression if a and b are both positive the fractionation refers to the transition state. Reactant fractionation can be considered by either a or b being negative.FIG.1.-Typical solvent isotope curve plotted as In (k,/k,) against x where x is the atom fraction of deuterium. The curvature parameter y measures how far the observed curve deviates at x = 3 from the straight line A. To carry out the inventory values of ay are plotted on a diagram of the type shown in fig. 2. Each parabola is labelled with the ratio a/b and where the line representing ay crosses a parabola we have a possible transition state. For instance in fig. 2 where the line representing y crosses the parabola labelled 1we have a = b = I. The parabola labelled 0 represents a model with the 4Aprotons singled out and a differential medium effect (b 3 a),where lots of & protons each contribute a little. The position of the intersection allows one to read off the values of AA and A,.These parameters describe the relative contributions of the 4Aand 4Bprotons to the overall effect and from these values of 4Aand 4Bcan be calculated. The two particular transition states discussed by Schowen are represented by S for a single proton in flight with no contribution from cbs and by Q for two protons contributing W. J. Albery Proton Transfer Reactions (Chapman and Hall London 1975) p. 272. 160 GENERAL DISCUSSION equally. From the data in Schowen’s paper we obtain the values of y in table 1. It can be seen that for systems (3) and (5) no conclusions can be drawn. For system (4) and probably less certainly system (2) the single proton model is correct. System (1) is plotted in fig.2 and the solvent isotope effect does not allow us to decide between FIG.2.-Plot of y for chymotrypsin acylation reaction to find possible transition states with fraction- ation of the form (1 -x+ $AX)(1 -x+ $BX)~. The parabolas are labelled with a/b = 6-1for a = 1. Each intersection corresponds to a possible transition state and AA and AB describe the relative importance of the $A and $B factors to the overall solvent isotope effect. The broken lines are one standard deviation for the value of y. An intersection at S corresponds to a transition state with a single proton in flght and no other significant factors ; an intersection at Q corresponds to a transition state with two protons in flight with +A = $B. the single or the quadratic model. Systems (6) and (7) have low values of y which arise from the “quadratic ” shape discussed by Schowen.However values of the fractionation factors for different possible transition states for system (6) are collected in table 2. The models with b = 10 correspond to a medium effect. Although TABLE 1 .-VALUES OF y no. system chymotrypsin acylation 0.65+ 0.13 chymotrypsin deacetylation 1.3 k0.3 trypsin 1.4 kO.9 elastase 1.10_+0.04 glutaminase -0.3 kO.9 asparaginase E. Coli 0.31+0.09 asparaginase Erwinia 0.36f 0.15 TABLE 2.-vALUES OF FRACTIONATION FACTORS FOR POSSIBLE TRANSITION STATES FOR SYSTEM 6 a 1 1 2 2 4~ 0.60 0.54 0.66 0.65 b 3 10 3 10 $B 0.82 0.95 0.92 0.98 this effect or the secondary effect for b = 3 are closer to unity for the model with two protons in flight (a = 2) the models with one proton in flight (a = 1) cannot be completely ruled out.Furthermore the values of 4Afor the two proton model are rather large for protons in flight. Remembering that for the fractionation in L30+ 162 GENERAL DISCUSSION 1 = 0.69 it might be that the low values of y for systems 6 and 7 are caused by changing -NH2 to -N+H3 and there are no protons in flight at all. Prof. R. L. Schowen (University of Kansas) said Albery’s method of representing acceptable mechanistic models for interpreting rates in mixed light and heavy water is very admirable for its simplicity clarity and economy. It should be particularly useful in cases where only a few data are available. Our customary method of treating such results is more conventional.We first fit the data for k to a polynomial in n by least-squares polynomial regression evaluating by Fisher’s F test the statistical significance of each term in the polynomial equation. This specifies the number of “active” protons which are required by the data. Thus if only the linear term is significant a one-proton system is indicated if the quadratic term is significant two protons are required etc. Having thus found the number of protons required and thus the number of factors needed on the r.h.s. of eqn (4b) of our paper we then fit k to the appropriate form of this equation by a general least-squares technique. This provides the best-fit values of the fractionation factors. Other things being equal one should obtain similar results from Albery’s method and the method just described.ERWlNlA L -ASNASE (V,/V,) = 2.52( I-n+n/ 1.62) SL’BSTFIPTE I-ASN 0 0.5 1 n FIG.1.-Ratio of maximum velocity V in mixed isotopic solvent to V for deutzrium oxide as a function of n for hydrolysis of asparagine (upper curve) and glutamine (lower curve) by asparaginase of Erwiniu curotovuru. The conditions are the same as for fig. 6 of our paper. If we use the numbering for the different systems given in Albery’s table 1 our treatment finds the linear term for all systems (1 to 7 inclusive) significant at the 0.999 level and the cubic term not significant above the 0.8 level (systems 2 and 4) or the 0.9 level (systems 1 3 5 6 and 7). This shows either one or two protons to be GENERAL DISCUSSION 163 required by all the systems studied.For systems 2 and 3 the quadratic term is not significant at the 0.8 level confirming the conclusion of one-proton catalysis. If one attempts nevertheless to fit a two-proton form of eqn (4b) to the data for system 2 for example the fractionation factors found correspond to isotope effects of 2.53 and 0.96 for the two protons further showing that only one proton is effectively responsible for the solvent isotope effect. For system 1 the quadratic term is significant at the 0.95 level and for system 4 at the 0.99 level. However the best-fit fractionation factors for the two-proton model yield isotope effects of 1.69 and 1.14 (system 1) and 2.50 and 0.98 (system 4). Thus once again it is a single proton in effect which produces the entire solvent isotope effect.For the amidohydrolase systems (5 to 7 inclusive) the quadratic term is significant at the 0.999 level in all cases. The best-fit fractionation factors are in reasonable agreement with the square- root dependencies shown in our paper. For these reactions therefore two-proton catalysis is indicated. Thus our usual method of data treatment indicates that for systems 2 3 and 4 no contribution beyond a few percent is to be ascribed to a second proton while system 1 can tolerate a second-proton isotope effect of about 1.14 (versus 1.69 for the first proton). For the amidohydrolases (systems 5-7) two-proton catalysis is fully confirmed. The discrepancies between these conclusions and Albery’s perhaps arise from the use of the entire data set in the method of treatment just discussed.The models discussed by Albery with 3-10 reactant-state contributing protons are closely related to models used earlier by Kresge and are best discussed below in connection with Kresge’s remarks. It can however be noted here that the involvement of ammonium functional groups is unlikely to product unusual effects \ \ \ since the fractionation factors for /N-H and -N-N+ both appear to be unity.’ / In the period following submission of our paper some support for the view expressed in its last sentence has been obtained by Mr. Daniel M. Quinn Mr. Mark Patterson and Mr. Robert Jarvis. If the two-proton catalytic entity of the amido- hydrolases exhibits coupled proton motion because its overall length in the catalytic transition state is just correct for the natural substrates studied (glutamine with glutaminase and asparigine with aspariginases) this length might be altered and the coupling destroyed if unnatural substrates were employed.Fig. 1 shows data for Erwinia asparaginase with the natural substrate asparagine (showing two-proton catalysis) and the unnatural substrate glutamine. For the latter (which reacts more slowly by a factor of about 30) the solvent isotope effect is reduced and VJn)becomes linear. Thus alteration of the substrate structure converts Erwinia asparaginase from a two-proton catalyst (with asparagine) to a one-proton catalyst (with glutamine). It is notable that only a very modest acceleration factor (maximally 30-fold) is associated with the coupling.Dr. W. J. Albery (Oxford) (communicated) First let me emphasise that in calculating the value of y I did use the complete data set for each system. Second I would like to comment on Schowen’s statement that the cubic term is insignificant and “ this shows either one or two protons to be required by all the systems studied ”. It is tempting to conclude that a product of the form n(1-x+4,x). . . (I -x+4,x) I R.L.Schowen. Progr. Phys. Org. Chem.. 1972 9 275. GENERAL DISCUSSION will yield a significant term in x“. But alas this is not the case. This can be demon- strated by considering the following particular example 53 15 (1-x++43 = 1-54~2+~ where A = X( 1 -X)(X -9/27.Now for all values of x (0 < x Q l) 1A1 is smaller than 2 x and thus compared to the normal experimental scatter A is insignificant. Hence it is not surprising that in fitting a solvent isotope curve to a polynomial the cubic term is insignificant. This will be true whatever the number of terms in the fractionation product. All solvent isotope curves can be satisfactorily fitted with just the linear and the quadratic terms. Thus one cannot find the number of “active protons” by evaluating the statistical significance of terms in a polynomial. This is the reason why the extra information from measurements in H20/D,0 mixtures can be described by the single curvature parameter y. It is also the reason why it is better to concentrate the mixture measurements at x = 3.Returning to Schowen’s systems the two different methods of analysis agree about systems 2 and 4. With respect to systems 3 and 5 the y treatment shows that the experimental scatter is too large to discriminate between the different models; it is true that the single proton model fits system 3 better than the two proton model and vice versa for system 5 but the alternative models cannot be ruled out. Incidentally one advantage of the y treatment is that one can imagine a Gaussian curve constructed on the y plot and this allows one to visualise how tentative or otherwise one’s con- clusions must be. Systems 6 and 7 are similar to each other. Table 1 shows for TABLE 1.-VALUES OF (kx/kl,-,bs)FOR DIFFERENT TRANSITION STATES^ FOR SYSTEM 6 a 1 1 2 2 3 0.602 0.543 0.658 0.649 0.702 4A b 3 10 3 10 3 0.821 0.952 0.917 0.977 0.987 4B T obs 0.00 2.92 2.93 2.93 2.93 2.93 2.93 0.10 2.66 2.67 2.67 2.67 2.67 2.67 0.20 2.47 2.42 2.42 2.42 2.42 2.42 0.30 2.18 2.19 2.19 2.19 2.19 2.19 0.40 1.97 1.97 1.98 1.97 1.97 1.97 0.50 1.77 1.77 1.78 1.77 1.77 1.77 0.60 1.58 1.59 1.59 1.59 1.59 1.59 0.70 1.42 1.42 1.42 1.42 1.42 1.41 0.80 1.27 1.26 1.26 1.26 1.26 1.25 0.90 1.07 1.11 1.11 1.11 1.11 1.11 1.00 1.00 0.98 0.98 0.98 0.98 0.97 S.D.0.023 0.024 0.023 0.023 0.023 a The transition state fractionation is (I -x+ $AX)’(l -x+ $BX)~; b S.D. is the standard deviation. system 6 the fit obtained by the y treatment for the different transition states.As always the different models give the same solvent isotope curve. (The models with b = 3 and b = 10 do not describe reactant fractionation but transition state fractiona-tion). In considering models of this type one is not suggesting any chance cancelling of factors. On the contrary the models with b = 10 correspond to a medium effect GENERAL DISCUSSION 165 or to small contributions from a large number of sites which do not cancel. The data in table 1 confirm the earlier conclusion that as good a fit to the data can be found for the single proton model as for the double proton model. I have conducted a similar analysis of system 1. Possible models with associated fractionation factors are given in table 2.Again all the different models give virtually the same solvent TABLE 2.-POSSIBLE TRANSITION STATES FOR SYSTEM 1 U 1 1 1 2 2 4A 0.608 0.599 0.597 0.695 0.694 b 1 3 10 3 10 $I3 0.868 0.958 0.988 1.030 1.009 103(s.~.) 9.3 9.3 9.4 9.2 9.2 isotope curve and for each model I have given in table 2 the standard deviation between the calculated curve and the experimental curve. The model in the left-hand column (a = 1 b = 1) is the same model found by Schowen (4A1 = 1.64 and & = 1.15) ; however as the two right-hand columns show models with two protons and a very modest medium effect will fit the data equally well. Finally the results given by Schowen in his last paragraph are most interesting and suggest that the alteration of y with substrate may be the best way to investigate the two proton mechanism.Prof. R.L. Schowen (University of Kansas) (communicated):Polynomial regression is capable of accurately measuring the number of active protons when the requisite precision is available and it always indicates the number of active protons required by the data. Models having more active protons cannot then derive support from the data (see below). The precision needed to establish a given term is a sensitive function of the isotope effect. In Albery’s example with 4 = 0.67 the third proton requires precision of a few tenths of a percent and the second proton about three percent. At 4 = 0.57 one percent suffices for the third proton and at 4 = 0.47 two percent. Albery has also given further detail on many-proton fits to system 6 (numbering of Albery’s original table l) for which polynomial regression shows linear (confidence level 0.999) and quadratic (0.999) terms to be highly significant and the cubic term (0.9) much less so.We accordingly interpret the results in terms of two-proton catalysis. Albery shows in his new table 1 that there exist many proton models according to which he can split the contribution of the second proton into three parts (or equally well into lo) with a standard deviation of 1-2 % or that he can pare away factors of up to around 1.2 and split these into numerous parts or that he can build on the less significant cubic term. Again I will have to refer to my reply to Kresge there seems to me little profit in constructing enormously complex models on the basis of data which justify only the simple two-proton picture.Similar comments apply to system 1 (Albery’s new table 2) interpreted by us as a one-proton case because of the low significance of the quadratic (0.95) and cubic (0.8) terms in the polynomial. As Albery has noted the two-proton model in the leftmost column of this table is that to which I referred above in saying “ . . it is a single proton in effect which produces the entire solvent isotope effect ”. At the right end of this table are examples in which “ complex cancellation of contributions . . . so chosen as not to perturb the linearity of v,(n) to an easily detectable extent” are again adduced as previously by Kresge.This Symposium page 160. GENERAL DISCUSSION As I mentioned before the y-method and polynomial regression ought to give similar results and I speculated that use of the whole data set in polynomial regression might explain the discrepancies. From the more detailed presentation I now see that the whole data set is used in both methods. The apparent discrepancies I now suppose to correspond to the allowance by the y-method of models which would produce terms with low confidence limits (say less than 0.8-0.9) in the polynomial regression. These are indeed possible models which greater precision might later justify. Finally I completely agree with the burden of Albery's last comment that investigation of structural and other effects on the shapes of solvent isotope effect curves represents the most promising avenue of pursuit of the questions in this field.Prof. A. J. Kresge (Urziuersityof Toronto)said It is well known that enzymes have many labile hydrogens which exchange rapidly with an aqueous solvent and con- formational and other changes in the enzyme that occur during catalysis could produce secondary isotope effects at these labelled positions which might combine to mark the isotope effect due to proton transfer.' To investigate this matter we have begun measuring solvent isotope effects on certain reactions of cycloamyloses. These substances are cyclic glucopyranosides with central cavities-" active sites "-which mimic the catalytic action of some enzymes remarkably closely ; they are at the same time relatively small molecules (6-8 glucose residues) with a manageable number of exchangeable sites (3 per glucose residue).We have found so far that dissociation of the cyclohexylamylose-p-nitrophenylate inclusion complex gives an overall solvent isotope effect of 0.72 (KD/KH) and an effect of 0.91 (Kx/KH)in a 50 50 H20-D20 mixture (X= 3). This curved depend- ence of isotope effect on solvent deuterium content has just the form needed to convert the behaviour expected for two-proton transfer catalysis into that for a one-proton transfer mechanism. Prof. R.L. Schowen (Uniuersity of Kansas) said Kresge's interesting investigation confirms that small solvent isotope effects may be generated by molecular association processes. This is to be expected and one can hope that complete elucidation of these effects will lead to an improved understanding of the detailed structural changes which accompany association processes including the association of substrates with enzymes.However I do not consider that these observations suggest any changes in the interpretation of our results for serine proteases. In the paper to which he has referred Kresge proposed an alternative interpreta- tion of our data for system 2 of Albery's table 1. Whereas we ascribed the linear character of v,(n) to the generation of the entire solvent isotope effect by a single transition-state proton Kresge proposed models in which 75-80 % of the solvent isotope effect arose from a single transition-state proton and the remainder of the effect from a complex cancellation of contributions of up to 21 other reactant and transition-state protons.These contributions were so chosen as not to perturb the linearity of u,(n) to an easily detectable extent. As is explained in our paper such a complex cancellation might conceivably account for a single case of linear u,(n) but to believe that a similar cancellation (involving different magnitudes and different physical processes) should hold for various substrates enzymes and rate-determining steps would require an extraordinary act of faith. It also happens that none of the A. J. Kresge J. Amer. Chem. SOC.,1973 95 3065. D. W. Griaths and M.L.Bender Adu. Catalysis 1973 23,209. GENERAL DISCUSSION 167 cases for which we report a linear v,(n) (all being maximum-velocity studies) involves a molecular association process.Since the submission of our paper we have completed a few further investigations of serine proteases. For example linear zI,(n) with linear term significant at 0.999 level is observed for deacylation of a-N-benzoyl-L-arginyltrypsin(quadratic and cubic terms not significant at level of 0.8) and a-N-benzoyl-L-arginylthrombin (quadratic term significant only at 0.9 cubic not significant at 0.8 level) and for acylation of elastase by a-N-carbobenzyloxy-L-alanine p-nitrophenyl ester (quadratic term not significant at level of 0.8 cubic only at 0.8). The former two cases do not involve association processes while the last case does. In my opinion to pursue a model in which exact cancellation between reactant- state and transition-state contributions occurs in each of these cases is quite unlikely to prove fruitful.The current finding of one-proton catalysis does not of course mean that if the physiological substrate structure is approached more closely for the serine proteases no coupling will be observed. Indeed the amidohydrolase results suggest that coupling could occur with proper choice of substrate. If so we may expect to observe non-linear v,(n) with the serine proteases. Prof. R. P. Bell (University of Stirling) said Schowen has made theoretical calculations of the relation between the lengths of hydrogen bonds and the occurrence of coupling between the motion of two protons and has used the results to draw conclusions about the nature of such motion in enzyme reactions.I would like to ask whether this relation has been tested in model systems with known hydrogen bond lengths or whether such tests are feasible. Prof. R. L. Schowen (University of Kansas) said The only attempt of which I know to simulate the charge-relay system in a small molecule is the effort of Rogers and Bruice which has been discussed together with other pertinent information by Bruice.2 The hydrogen-bond lengths in this molecule are not known. Although there should in-principle be no barrier to the construction of suitable model com- pounds it may emerge that enzymes (for which structures are now rapidly becoming available) offer the best field for testing the calculations.Dr. R. A. More O’Ferrall (University College Dublin) said It was first shown some years ago in the general base catalysed cyclisation of chlorobutanol that ACO-H-o-c-c1 proton transfers between oxygen atoms at least when concerted with bond-making or bond-breaking to carbon may be characterised by primary hydrogen isotope effects abnormally close to unity. I am not sure that this finding has been fully explained but I wonder if it is relevant to the interpretation of Knowles and Albery’s result ? G. A. Rogers and T. C. Bruice J. Amer. Chem. SOC.,1974 96,2473. T. C. Bruice Ann. Rev. Biochem. 1974. C. G. Swain D. A. Kuhn and R. L. Schoweo J. Amer. Chem. SOC.,1965,87 1553. GENERAL DISCUSSION Prof. R. L. Schowen (Uniuersity of Kansas) said The conclusion reached in the paper which More O’Ferrall has cited is that the proton is not “ in flight ” in the transition state for reorganization of the heavy-atoms ; that is it has no significant amplitude in the reaction coordinate.I still hold this view and suspect that it applies both to the kind of case observed by Fisher Albery and Knowles in which the fractionation factor for the critical proton is near unity and to a large number of cases in which the fractionation factor is around 0.5. Even if this is so no inter- mediate structure (in the usual sense of a stable species) is necessarily required between the transition state and the product in which the proton has been transferred. Instead as Choi and Thornton have explained the reaction path from the heavy- atom reorganization transition state may lead directly down into the proton-transfer transition state entering at a right angle to the proton-transfer reaction path.Dr. W. J. AIbery (Oxford University) and Prof. J. R. Knowles (Harvard University) (communicated) Although the fractionation factor for a proton transfer that is concerted with other covalency changes may be close to unity in our case the proton transfer from the carboxylate base would be a simple single step. Thus the factor for this type of transfer one would still expect to be less than 0.3. Prof. R. A. Marcus (University of Illinois) said Did you consider the possibility of investigating the effect of added molecules that might more readily accept a proton? Prof.J. R. Knowles (Harvard University) said No. I have considerable doubts about doing enzyme-catalyzed reactions in anything but buffered solutions in IH2O. I am even prejudiced against 2H20(pace Professor Schowen) since the effects upon such things as the apparent pK,-values of ionising groups on the enzyme the con- formation of the proton and effects upon the exact constellation of active-site function- alities are largely unknown and are very hard to assess. [Professor W. P. Jencks communicated the possibility that the proton transfer could be dependent on the buffer (in our case triethanolamine). Our data cannot rule this out though the steric constraints of the enzyme’s active site make direct transfer to buffer unlikely even though Grotthus-type transfer via the solvent is undoubtedly possible.] M. Choi and E. R. Thornton J. Amer. Chem. Suc. 1974 96 1428.

ISSN:0301-5696    

DOI:10.1039/FS9751000160

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

AUTHOR INDEX* Albery W. J. 90 95,98 132 137 138 139 142 154 160 163 168. Arnett E. M. 20 50 51. Banger J. 113. Bannister C. E. 78. Bell R. P. 7 51 91 137 139 167. Bordwell F. G. 51 100 132 133. Caldin E. F. 51 53 97 121 138 140. cox B. G. 107 134 136. Crooks J. E. 29 52 54. Eggins B. R. 134. Elrod J. P. 145. Fendler. J. H. 57. Fisher L. M. 154. Gandour R. D. 145. Gibson A. 107. Goodall D. M. 94. Hassid A. I. 69. Hogg J. L. 145. Hyne J. B. 135. Jaffe A. 113. Jencks W. P. 41 58. Johnston D. E. 20. Jones J. R. 50. Kise M. 145. Knowles J. R. 154 168. Kreevoy M. M. 55 69 92 97 99. Kresge A. J. 89 133 166. Liang T.-M. 69. Ein AX. 113. Maggiora G. M. 145. Marcus R. A. 60 91 93 94 141 168. Margerum D. W. 78 98 99. Melander L. 52 91. More O’Ferrall R. A, 58 92 132 167. Oancea D. 20. Parbhoo D. M. 55. Raychera J. M. T. 78. Robinson B. H. 29 52 54 55. Saunders W. H. 50 58 97 113 136 137 138. Sayer J. M. 41. Schowen R. L. 145 162 165 166 167 168. Small L. E. 20. Venkatasubban K. S. 145. Wilson C. J. 121. Wong L. F. 78. Zuman P. 138. * The references in heavy type indicate papers submitted for discussion.

ISSN:0301-5696    

DOI:10.1039/FS9751000169

出版商:RSC    

年代:1975

数据来源: RSC